MEMCAL    SCHOOL 


COLLEGE  OF  ESARHACT 


CALIFORNIA    COLLEGE 
gf  PHARMACY 


:    .      •">• 


««cfl»  of  Pharmacy 

- 


A  LABORATORY  MANUAL 


OF 


QUALITATIVE 
CHEMICAL  ANALYSIS 

FOR  STUDENTS  OF  PHARMACY 


BY 


THEODORE  J.  BRADLEY,  A.M.,  B.S.,  Ph.G. 

DEAN  AND  PROFESSOR  OP   ANALYTICAL,   AND   ORGANIC  CHEMISTRY 
IN  THE  MASSACHUSETTS  COLLEGE  OF  PHARMACY 


California  Colte^e  of  Pharmacy 


LEA   &   FEBIGER 

PHILADELPHIA   AND    NEW    YORK 
1920 


COPYRIGHT,   1917 
BY  THEODORE  J.  BRADLEY 

COPYRIGHT,  1920 
BY  LEA  &  FEBIGER 


The  use  in  this  volume  of  certain  portions  of  the  text  of  the  United  States 
Pharmacopoeia  is  by  virtue  of  permission  received  from  the  Board  of  Trustees 
of  the  United  States  Pharmacopoeial  Conventiqn.  The  said  Board  of  Trustees 
is  not  responsible  for  any  inaccuracy  nor  for  any  errors  in  the  statement  of 
quantities  or  percentage  strengths. 


PREFACE. 


QUALITATIVE  analysis  is  of  use  to  the  pharmacist  in  testing 
chemicals  for  identity  and  purity.  The  study  of  the  subject  is 
also  of  value  because  of  the  practical  knowledge  of  chemicals  and 
chemical  processes  acquired  by  the  student,  and  because  it  gives 
training  in  careful  observation  and  develops  the  reasoning  powers 
as  they  are  exercised  in  the  interpretation  of  results. 

This  Manual  was  prepared  as  a  guide  for  the  author's  own  classes, 
and  the  objects  sought  are  to  acquaint  the  student  with  the  general 
methods  of  qualitative  analysis  and  to  prepare  him  to  carry  out  such 
qualitative  tests  as  the  pharmacist  may  be  called  upon  to  make. 
The  course  is  arranged  to  include  one  hour  of  lecture,  one  hour  of 
recitation  and  about  three  hours  of  laboratory  work  per  week  for 
one  school  year.  Practice  on  the  analysis  of  unknown  solutions 
is  provided  for  throughout  the  course.  This  is  important,  as  it  not 
only  increases  the  interest  of  the  student  in  the  work  but  also 
develops  his  self-reliance  by  constantly  putting  him  upon  his  own 
responsibility  in  doing  his  work  and  in  interpreting  his  results. 
The  introductory  section  on  the  theory  of  chemistry  may  be  omitted 
if  it  is  not  necessary  for  the  class  to  study  or  review  this  part  of  the 
subject. 

The  Manual  is  in  no  sense  a  reference  book,  and  in  general  only 
those  things  are  included  that  are  needed  as  a  guide  for  the  labora- 
tory work,  or  which  may  easily  be  learned  in  connection  with  this 
work.  A  standard  reference  book  on  qualitative  analysis  will  be  of 
great  service  while  taking  the  course,  and  such  a  reference  book  is 
indispensable  for  anyone  who  is  to  pursue  the  subject  further. 

T.  J.  B. 

BOSTON,  MASS.,  1920. 


CONTENTS. 


ELEMENTARY  THEORY  OF  CHEMISTRY. 

Matter 9 

Changes  in  matter 9 

Laws  of  definite  and  multiple  proportion 10 

Structure  of  matter 11 

Composition  of  matter 11 

Kinds  of  chemical  changes 12 

Symbols 13 

Formulas 13 

Chemical  equations 13 

Atomic  weight • 14 

Valence 14 

Acids,  bases  and  salts 15 

Compound  radicals 16 

Nomenclature 17 

Solution 19 

Water  of  crystallization         20 

Chemical  analysis 21 


THE  METALS. 

Groups  of  the  metals 22 

Group  1 23 

Lead 23 

Mercurous  mercury 25 

Silver 26 

Analysis  for  group  1 27 

Group  2 28 

Mercuric  mercury 29 

Bismuth 30 

Copper 31 

Cadmium 33 

Analysis  for  group  2 34 

Group  3 37 

Gold    . 37 

Platinum 38 

Arsenic     .  39 

Antimony 41 

Tin 43 

Analysis  for  group  3 44 


vi  ,  CONTENTS 

Group  4 48 

Iron 48 

Chromium 51 

Aluminum                        .    • 52 

Analysis  for  group  4 53 

Group  5 57 

Cobalt 57 

Nickel 58 

Manganese 59 

Zinc 61 

Analysis  for  group  5 63 

Group  6 67 

Calcium 67 

Strontium 68 

Barium 69 

Analysis  for  group  6 70 

Group  7 . 74 

Magnesium 74 

Potassium 76 

Sodium 77 

Lithium 79 

Ammonium 79 

Analysis  for  group  7 '  .      .  81 

Analysis  of  a  solution  for  all  groups  of  the  metals 82 


THE  ACIDS. 

Groups  of  the  acids 86 

Group  A  of  the  acids 86 

Hydrochloric  acid 86 

Hydrobromic  acid 88 

Hydriodic  acid 89 

Hydrocyanic  acid 90 

Hydrosulphuric  acid 91 

Analysis  for  group  A 92 

Group  B  of  the  acids       .      .      .      / 93 

Sulphuric  acid 93 

Sulphurous  acid 95 

Carbonic  acid 96 

Analysis  for  group  Bl    . 97 

Oxalic  acid 97 

Phosphoric  acid 98 

Analysis  for  group  B2 99 


CONTENTS  vii 

Group  C  of  the  acids 100 

Boric  acid 100 

Acetic  acid 101 

Nitric  acid 102 

Analysis  for  group  C .  103 

Analysis  of  a  solution  for  the  important  metals  and  acids 104 

Treatment  of  solid  substances 106 

Qualitative  examination  of  official  inorganic  chemicals 109 

General  qualitative  tests  of  the  United  States  Pharmacopoeia       .      .      .  115 

Destruction  of  organic  matter 119 

Reagents  and  test  solutions 120 

Index  127 


CALIFORNIA    COUtGE 
of   PHARMACY 


ELEMENTARY  THEORY  OF  CHEMISTRY. 


A  part  or  all  of  this  section  may  be  omitted  if  it  is  not  necessary  for  the  class  to 
study  or  review  this  subject. 


Matter  is  anything  which  occupies  space,  as  wood,  water  and  air. 
In  addition  to  its  occupation  of  space,  which  is  called  its  extension, 
matter  has  several  other  general  properties,  including  gravitation, 
indestructibility,  inertia,  impenetrability  and  divisibility,  which  are 
studied  in  physics.  Some  of  the  things  which  exist  but  are  not 
material,  are  energy  in  its  various  forms,  mental  processes  and  the 
emotions. 

A  substance  is  a  separate  kind  of  matter,  like  iron,  marble,  water  or 
air.  Each  substance  possesses  all  of  the  general  properties  of  matter 
and,  also,  several  properties  which  serve  to  distinguish  it  from  other 
kinds  of  matter  and  are  called  specific  properties.  Among  these  are 
its  condition  of  being  a  solid,  a 'liquid  or  a  gas,  its  density,  and  some 
of  the  following:  color,  odor,  taste,  melting-point,  boiling-point, 
crystalline  form,  hardness,  brittleness,  tenacity,  behavior  toward 
other  substances,  and  some  other  properties.  The  particular  com- 
bination of  these  properties  possessed  by  a  substance  serve  to 
identify  it.  Most  substances  may  be  separated  into  several  other 
substances,  with  more  or  less  difficulty,  and  a  substance  is  said  to 
be  composed  of  its  constituents,  which  are  referred  to,  collectively, 
as  its  composition. 

Changes  in  Matter. — Matter  is  constantly  undergoing  change.  The 
changes  may  be  rapid  or  slow  but  nothing  is  permanent.  The 
changes  are  very  various,  but  they  may  all  be  classified  into  two 
great  kinds :  one  kind  affecting  the  condition  of  the  matter  and  the 
other  kind  affecting  its  composition.  The  changes  which  affect 
the  condition  of  matter  are  physical  changes  and  include  all  forms  of 
motion,  temperature  changes,  magnetic  changes,  changes  of  state, 
like  the  freezing  or  boiling  of  water,  and  many  others.  In  a  physical 
change  the  identity  of  the  matter  undergoing  the  change  is  not 
affected.  The  changes  which  affect  the  composition  of  matter  are 
called  chemical  changes  and  include  all  burnings  or  combustions, 
the  decay  of  vegetable  and  animal  matter,  the  rusting  of  iron,  and 


10  ELEMENTARY  THEORY  OF  CHEMISTRY 

various  others.  A  chemical  change  is  frequently  called  a  chemical 
reaction  or  simply  a  reaction.  In  a  chemical  change  the  identities 
of  the  matter  undergoing  the  change  are  lost  and  new  substances 
with  new  identities  are  formed. 

Physical  and  chemical  changes  are  often  mutually  dependent 
upon  each  other;  heat,  light,  electricity,  solution  and  other  physical 
agencies,  often  produce  chemical  changes  and,  conversely,  chemi- 
cal changes  often  produce  heat,  light,  electricity  and  other  physical 
effects.  Physical  and  chemical  changes  are  all  caused  by  energy, 
which  is  that  which  can  produce  change  in  matter.  Energy  cannot 
be  created  or  destroyed,  which  fact  is  referred  to  as  the  conservation 
of  energy.  Heat,  light,  electricity,  and  the  power  due  to  the  motion 
or  the  position  of  matter,  are  different  forms  of  energy,  and  these 
agencies  may  produce  chemical  change  or  be  produced  by  chemical 
change.  Matter  can  neither  be  created  nor  destroyed,  which  fact 
is  referred  to  as  indestructibility  or  the  conservation  of  matter. 

Physical  science  studies  the  properties  of  matter  and  the  effects 
of  energy  on  matter.  Physics  is  the  branch  of  physical  science 
which  studies  the  general  properties  of  matter  and  the  changes 
which  matter  undergoes  without  affecting  its  composition.  Chemis- 
try is  the  branch  of  physical  science  which  studies  the  specific 
properties  and  composition  of  substances  and  changes  in  their 
composition. 

Laws  of  Definite  and  Multiple  Proportions. — Early  in  the  develop- 
ment of  the  science  of  chemistry  two  important  laws  were  formulated 
from  the  results  of  experiments.  The  first  is  called  the  law  of 
definite  proportions  and  may  be  stated  as  follows:  In  any  chemical 
change  there  is  always  a  definite  relationship  between  the  amounts 
of  the  substances  entering  into  and  produced  by  the  change.  For 
example,  copper  and  sulphur,  when  heated  together,  will  react  upon 
each  other,  in  the  proportion  of  63  parts  of  copper  to  32  parts  of 
sulphur,  any  excess  of  either  substance  being  left  unchanged.  Also, 
there  is  a  definite  amount  of  the  resulting  substance  formed  by  the 
reaction  of  any  given  amount  of  copper  upon  the  corresponding 
amount  of  sulphur. 

The  second  law  is  called  the  law  of  multiple  proportions  and  may 
be  stated  as  follows:  If  two  substances  react  upon  each  other  in 
different  proportions  by  weight  the  amounts  of  each  substance  in. the 
different  reactions  bear  a  simple  relationship  to  each  other.  For 
example,  sulphur  will  react  upon  oxygen  in  two  ways:  in  one  of  the 
reactions  a  given  amount  of  sulphur,  as  32  parts,  reacts  upon  32  parts 
of  oxygen,  and  in  the  other  32  parts  of  sulphur  reacts  upon  48  parts 
of  oxygen.  Ln  these  two  different  reactions,  the  amounts  of  the 
sulphur  being  equal,  the  relative  amounts  of  the  oxygen  are  to  each 
other  as  two  to  three,  a  simple  relationship. 


COMPOSITION  OF  MATTER  11 

Structure  of  Matter. — A  mass,  or  body,  is  a  separate  portion  of 
matter,  like  an  iron  rod,  a  block  of  marble,  the  water  in  a  reservoir, 
or  the  air  in  a  room.  Any  appreciable  mass  may  be  divided  into 
smaller  masses,  but  it  appears  that,  if  this  division  were  carried  far 
enough,  extremely  small  particles  would  be  obtained  which  could 
not  be  further  subdivided  without  changing  the  identity  of  the 
substance.  These  particles  are  called  molecules.  They  are  so  small 
in  size  that  many  millions  are  contained  in  even  a  small  mass  of 
matter,  like  a  grain  of  sand,  and  there  is  no  machinery  fine  enough 
to  completely  separate  a  mass  into  its  molecules,  nor  is  it  possible 
to  see  them  even  with  high  magnification.  The  specific  properties 
of  a  substance  are  due  to  the  properties  of  its  molecules  and  their 
influence  on  each  other. 

Molecules  may  be  temporarily  subdivided  by  other  than  mechani- 
cal processes,  but,  when  this  is  accomplished,  the  identity  of  the 
substance  changes.  The  smaller  particles  which  are  contained  in 
molecules  are  called  atoms.  Atoms  are  the  indivisible  particles  of 
which  all  matter  is  composed.  They  are  united  with  each  other  to 
form  molecules,  the  atoms  existing  separately  only  for  an  instant  as 
they  pass  from  one  molecule  to  another.  Any  collection  of  mole- 
cules is  a  mass.  The  attraction  of  molecules  for  each  other  is  called 
cohesion  or  adhesion.  The  attraction  which  causes  atoms  to  unite 
with  each  other  to  form  molecules  is  called  chemical  affinity,  chemical 
attraction,  or  chemism. 

The  description  above  of  the  structure  of  matter  is  a  brief  state- 
ment of  the  very  important  atomic  theory  which  is  the  foundation 
of  the  science  of  chemistry  and  has  many  applications  in  physics. 
While  all  parts  of  it  cannot  be  rigorously  proved,  it  is  the  only  com- 
plete explanation  of  the  facts  of  chemistry  that  has  been  formulated, 
and  it  is  now  generally  accepted  as  being  true. 

Composition  of  Matter. — If  the  molecules  of  a  substance  are  all 
alike  and  the  atoms  in  each  molecule  are  alike,  the  substance  is  of 
the  simplest  possible  composition  and  cannot  be  separated  into  any 
other  substances  by  any  known  means.  Such  a  simple  substance 
is  called  an  element.  About  eighty  elements  have  been  discovered, 
each  of  the  countless  number  of  substances  known  being  composed 
of  one  or  more  of  these  elements. 

If  the  molecules  forming  a  substance  are  all  exactly  alike,  but  each 
molecule  contains  different  kinds  of  atoms,  the  substance  is  called  a 
compound.  The  number  of  possible  compounds  formed  from  the 
eighty  elements  is  very  large,  just  as  the  number  of  words  formed 
from  the  twenty-six  letters  of  the  alphabet  appears  to  be  unlimited. 
Many  compounds  are  found  in  nature  and  many  others  are  manu- 
factured. It  is  not  possible  to  separate  the  constituents  of  a 
compound  from  each  other  by  physical  processes,  but  a  compound 


12  ELEMENTARY  THEORY  OF  CHEMISTRY 

may  be  separated  into  simpler  compounds  or  elements  by  chemical 
processes,  such  a  separation  being  called  a  chemical  decomposition. 
In  a  compound  the  constituents  which  compose  it  are  united  with 
each  other  in  a  fixed  and  definite  proportion  by  weight.  The 
compound  has  a  separate  identity,  and  properties  which  are  very 
different  from  those  of  its  constituents. 

If  a  substance  is  formed  of  different  kinds  of  molecules  it  is  called 
a  mixture.  In  a  mixture  the  different  constituents  are  merely 
commingled  with  each  other,  and  each  retains  its  own  identity  and 
properties,  and  the  properties  of  the  mixture  are  between  the 
properties  of  the  constituents.  It  is  generally  possible  to  separate 
the  constituents  of  a  mixture  from  each  other  by  mechanical  pro- 
cesses, because  of  their  different  solubilities,  melting-points  or  other 
specific  properties,  and  the  relative  amounts  of  the  constituents 
may  vary. 

Many  natural  substances,  like  granite  and  air,  are  mixtures, 
while  others,  like  water,  salt  and  marble,  are  compounds,  and  a  few, 
like  gold  and  sulphur,  are  elements.  An  absolutely  pure  compound 
or  element  is  very  unusual,  the  impurities  being  mixed  with  the 
compound  or  element. 

An  element  is  said  to  be  in  the  nascent  state  for  the  instant  during 
which  it  is  being  set  free  from  a  compound  and  before  its  atoms  have 
united  with  each  other  to  form  molecules.  When  the  atoms  of  an 
element  have  united  to  form  molecules,  a  certain  amount  of  energy 
is  necessary  to  overcome  their  attraction  for  each  other,  so  an  ele- 
ment is  more  active,  chemically,  in  the  nascent  state,  as  its  atoms 
then  exhibit  their  full  chemical  affinities. 

Kinds  of  Chemical  Changes. — Combination,  when  two  or  more 
substances  unite  with  each  other  chemically  to  produce  one  sub- 
stance, as  the  burning  of  hydrogen  in  oxygen  to  form  water. 

Decomposition,  when  one  substance  forms  two  or  more  substances 
as  the  result  of  a  chemical  change,  as  the  decomposition  of  water  by 
an  electric  current  to  form  hydrogen  and  oxygen. 

Double  decomposition,  when  two  substances  react  upon  each  other 
to  produce  two  other  substances  by  an  exchange  of  constituents. 
This  is  the  commonest  kind  of  chemical  change,  and  many  examples 
of  it  will  be  studied. 

There  are  many  changes  which  belong  to  more  than  one  of  these 
kinds  of  chemical  changes,  one,  called  substitution,  being  quite 
common.  In  this  kind  of  a  change  an  element  reacts  upon  a  com- 
pound in  such  a  way  that  a  part  of  the  compound  is  displaced,  the 
first  element  taking  its  place,  as  the  reaction  between  zinc  and 
sulphuric  acid,  in  which  the  zinc  replaces  the  hydrogen  of  the  acid, 
forming  zinc  sulphate  and  setting  hydrogen  free. 


CHEMICAL  EQUATIONS  13 

Symbols.— Early  in  the  development  of  the  science  of  chemistry  it 
was  found  convenient  to  designate  elements  by  signs.  At  first  these 
were  astronomical  signs  and  geometrical  figures,  but  as  the  number 
of  known  elements  increased  this  method  was  found  to  be  faulty  and 
the  present  symbols  came  into  use.  The  symbol  of  an  element  is 
formed  from  the  English  or  some  other  name  of  the  element  by  taking 
the  first  letter,  alone,  or  with  some  other  letter  in  the  name.  The 
first  letter  of  a  symbol  is  always  a  capital,  and  if  there  is  a  second 
letter,  it  is  always  a  small  letter.  When  symbols  first  came  into 
use  they  were  mere  abbreviations  of  the  names  of  the  elements, 
and  they  are  sometimes  used  in  this  way  at  present.  In  most  cases, 
however,  they  stand  for  one  atom  of  an  element.  If  several  atoms 
are  to  be  indicated  the  symbol  is  multiplied  by  methods  discussed 
in  the  next  paragraph. 

Formulas. — Chemical  formulas  are  used  to  represent  molecules. 
Formulas  are  made  by  writing  the  symbols  of  the  elements  in  the 
molecules,  one  after  the  other.  If  a  molecule  contains  only  one 
atom  of  a  certain  element  the  symbol  of  that  element  indicates  this; 
more  than  one  atom  of  any  element  in  a  molecule  is  indicated  by 
an  inferior  figure  written  after  the  symbol  of  that  element,  thus 
H2O  is  the  formula  for  water  and  tells  us  that  each  molecule  of 
water  contains  two  atoms  of  hydrogen  and  one  atom  of  oxygen; 
H2SO4  is  the  formula  for  sulphuric  acid,  whose  molecules  each  con- 
tain two  atoms  of  hydrogen,  one  atom  of  sulphur  and  four  atoms  of 
oxygen.  When  for  any  reason  a  part  of  a  molecule  consisting  of 
more  than  one  atom  is  considered  as  a  group  to  be  multiplied,  it  is 
enclosed  in  a  parenthesis,  which  is  followed  by  the  inferior  figure, 
as  in  ammonium  sulphate  (NH4)2SO4,  whose  molecule  contains  two 
NH4  groups,  combined  with  the  SO4  group.  A  formula  generally 
represents  one  molecule  of  a  compound.  More  than  one- molecule 
are  shown  by  a  line  figure  written  before  the  formula,  thus  two 
molecules  of  ammonium  sulphate  are  written  2(NH4)2SO4. 

Chemical  Equations. — When  a  chemical  change  occurs,  it  is  fre- 
quently possible  to  express  what  occurs  by  the  use  of  symbols  and 
formulas  written  in  the  form  of  an  equation.  Besides  being  of  great 
convenience,  these  equations  express  exact  relationships  between 
the  quantities  of  the  substances  entering  into  the  reaction  and 
produced  by  it;  thus  the  decomposition  of  water  is  expressed  2H2O  = 
2H2  +  O2,  which  tells  us  that  when  water  is  decomposed  two  mole- 
cules yield  two  molecules  of  hydrogen  and  one  molecule  of  oxygen. 
Actually  very  many  molecules  are  involved  when  any  appreciable 
amount  of  water  is  decomposed.  This  equation  is  sometimes 
written  H2O  =  2H  +  O,  but  the  first  form  is  preferable,  as  it  is 
the  simplest  way  of  writing  all  atoms  combined  to  form  molecules, 
which  they  do  as  soon  as  they  are  liberated  from  other  molecules. 


14  ELEMENTARY  THEORY  OF  CHEMISTRY 

Chemical  equations,  like  mathematical  equations,  must  be  balanced, 
that  is,  they  must  have  the  same  number  of  atoms,  of  each  kind 
involved,  on  both  sides  of  the  equality  sign. 

Atomic  Weight. — Atoms  are  so  extremely  small  that  their  absolute 
weights,  in  terms  of  any  ordinary  unit,  are  very  small  decimals 
which  have  been  calculated,  approximately,  but  they  cannot  be 
determined  directly.  The  relative  weights  of  the  atoms  of  different 
elements  have  been  quite  accurately  determined  and  have  been 
referred  to  the  weight  of  an  atom  of  hydrogen  as  a  unit.  What  is 
known  as  the  atomic  or  combining  weight  of  an  element  is  the  number 
of  times  that  the  weight  of  its  atom  is  heavier  than  the  weight  of 
an  atom  of  hydrogen.  Thus  the  atomic  weight  of  oxygen  is  15.88, 
which  means  that  an  atom  of  oxygen  is  15.88  times  as  heavy  as  an 
atom  of  hydrogen;  the  atomic  weight  of  iron  is  55.5,  which  means 
that  an  atom  of  iron  weighs  55.5  times  as  much  as  an  atom  of 
hydrogen,  and  so  on.  The  sum  of  the  weights  of  the  atoms  in  a 
molecule  of  a  compound  is  the  molecular  weight  of  that  com- 
pound. 

Valence. — The  atoms  of  different  elements  differ  from  each  other  in 
the  number  of  atoms  of  other  elements  that  they  can  combine  with 
chemically  to  form  molecules.  This  quantity  of  combining  power 
of  the  atoms  of  an  element  is  called  its  valence  and  is  measured  by 
the  number  of  atoms  of  hydrogen  that  one  atom  of  the  element  can 
combine  with,  or  take  the  place  of.  The  unit  of  valence  is  called 
a  bond.  An  element  whose  atom  combines  with,  or  displaces,  one 
atom  of  hydrogen,  as  Cl  in  HC1  and  Na  in  NaCl,  has  one  bond  and 
is  called  a  monad  element  or  simply  a  monad.  An  element  whose 
atom  combines  with,  or  displaces,  two  atoms  of  hydrogen,  as  O  in 
H20  and  Ca  in  CaO,  has  two  bonds  and  is  called  a  dyad.  Similarly 
a  triad  has  three  bonds;  a  tetrad,  four;  a  pentad,  five;  a  hexad,  six. 
There  are  very  few  compounds  in  which  an  element  appears  to  have 
a  valence  greater  than  six.  The  valence  of  an  element  is  indicated, 
when  necessary,  by  the  use  of  superior  Roman  numerals  written 
after  the  symbol  of  the  element,  as  H',  O",  Civ,  etc. 

Some  of  the  elements  have  the  same  valence  in  all  of  their  com- 
pounds, while  others  vary  in  valence,  forming  two  or  more  series 
of  compounds.  When  atoms  combine  to  form  molecules  all  bonds 
must  be  satisfied,  the  number  of  atoms  in  each  molecule  being  the 
smallest  number  that  will  comply  with  this  general  law.  If  two 
elements  have  the  same  valence,  one  atom  of  each  will  form  a  mole- 
cule of  their  compound,  as  two  monads  in  H'Cl',  Na'Cl',  two  dyads 
in  Fe"O",  Ca"S",  etc.  Two  atoms  of  a  monad  combine  with  one 
atom  of  a  dyad,  as  in  H'2O";  three  atoms  of  a  monad  with  one 
atom  of  a  triad,  as  in  N'^H'a;  two  atoms  of  a  dyad  with  one  atom 
of  a  tetrad,  as  in  C1VO"2;  two  atoms  of  a  triad  with  three  atoms 


ACIDS,  BASES  AND  SALTS  15 

of  a  dyad,  as  in  As'"2O"3;  two  atoms  of  a  pentad  with  five  atoms  of 
a  dyad,  as  in  PV2O"5,  and  so  on. 

Acids,  Bases  and  Salts. — Some  compounds  containing  hydrogen 
have  several  properties  in  common  and  are  called  acids.  They 
generally  have  a  sour  taste  and  change  certain  dyes,  like  litmus, 
from  blue  to  red.  Many  of  the  acids  contain  oxygen  combined  with 
the  hydrogen,  and  the  hydrogen  alone  or  the  hydrogen  and  oxygen 
are  combined  with  some  other  element,  called  the  characteristic 
element  of  the  acid.  Another  class  of  compounds,  called  bases, 
contain  oxygen,  or  oxygen  and  hydrogen,  combined  with  another 
element,  called  the  characteristic  element  of  the  base.  An  acid  will 
exchange  hydrogen  for  the  characteristic  element  of  a  base  form- 
ing water  and  a  compound  which  is  called  a  salt.  This  reaction 
between  an  acid  and  a  base  is  called  a  neutralization.  Examples: 
CaO  +  2HC1  =  CaCl2  +  H2O,  Ca(OH)2  +  2HC1  =  CaCl2  +  2H2O. 
Salts,  however,  may  be  formed  in  several  other  ways. 

An  element  which  combines  with  oxygen,  or  oxygen  and  hydrogen, 
to  form  a  base  is  called  a  metal  and  is  electro-positive.  An  element 
which  combines  with  hydrogen,  or  hydrogen  and  oxygen,  to  form  an 
acid  is  called  a  non-metal  and  is  electro-negative.  The  classification 
of  elements  into  metals  and  non-metals  has  long  been  used  and  was 
formerly  based  on  physical  properties.  It  is  not  exact,  as  there  are 
some  elements  which  behave  both  as  metals  and  non-metals  under 
different  conditions.  The  differences  are  in  degree,  some  elements 
always  forming  bases,  like  potassium  and  calcium,  some  always 
forming  acids,  like  chlorine  and  sulphur,  and  a  few  forming  both 
bases  and  acids,  like  antimony.  There  are  about  fifty  metals,  and 
they  vary  greatly  in  properties.  Some  are  familiar  substances,  like 
iron,  gold,  copper,  etc.,  while  others  are  seldom  seen  in  the  free  state 
and  can  only  be  obtained  with  difficulty  from  their  compounds. 
The  metals  are  opaque  -and  have  a  lustre  when  they  are  freshly 
cut  and  clean.  They  are  conductors  of  both  heat  and  electricity. 
Potassium,  sodium,  calcium,  magnesium,  zinc,  iron,  silver,  copper, 
mercury,  lead,  tin,  gold  and  aluminum  are  important  metals.  An 
alloy  is  formed  by  fusing  two  or  more  metals  together.  Some  alloys 
appear  to  be  unstable  compounds  and  others  appear  to  be  mixtures. 
There  are  many  important  and  widely  used  alloys,  such  as  brass, 
type-metal  and  solder. 

The  non-metals  vary  in  properties  even  more  than  the  metals  do. 
The  principal  non-metals  are  hydrogen,  oxygen,  sulphur,  chlorine, 
bromine,  iodine,  nitrogen,  phosphorus,  carbon,  silicon  and  boron. 

An  acid  containing  one  atom  of  replaceable  hydrogen  in  the 
molecule  is  said  to  be  monobasic;  similarly,  acids  with  two,  three  and 
four  atoms  of  replaceable  hydrogen  in  the  molecules  are  di-basic, 
tri-basic  and  tetra-basic,  respectively.  There  are  few  acids  with 


16  ELEMENTARY  THEORY  OF  CHEMISTRY 

more  than  four  replaceable  hydrogen  atoms  in  the  molecules.  An 
acid  composed  of  hydrogen  and  one  other  element  and  no  oxygen 
is  called  a  hydracid.  These  are  few  in  number:  hydrochloric  acid, 
HC1;  hydrobromic  acid,  HBr;  hydriodic  acid,  HI;  hydrofluoric  acid, 
HF;  and  hydrosulphuric  acid,  H2S,  comprising  the  entire  list  of 
hydracids.  Nearly  all  of  the  many  other  acids  contain  oxygen  in 
addition  to  the  hydrogen  and  the  characteristic  element  and  are 
called  oxy acids.  An  acid  anhydrid  is  a  compound  of  a  negative 
element  with  oxygen  which  will  combine  with  water  to  form  an  acid, 
or  which  is  formed,  along  with  water,  when  an  acid  is  decomposed, 
as  S03  combines  with  water  to  form  H2SO4;  and  CO2  is  formed, 
along  with  H2O,  on  the  decomposition  of  H2C03.  Different  acids 
may  be  obtained  from  the  same  negative  element,  if  it  combines  with 
different  amounts  of  oxygen  to  form  different  anhydrides,  or  one  of 
its  anhydrides  combines  with  different  amounts  of  water. 

A  base  may  be  either  a  compound  of  a  metal  and  oxygen  or  of  a 
metal,  oxygen  and  hydrogen.  A 'base  that  is  soluble  in  water  is 
called  an  alkali.  Alkalies  turn  red  litmus  back  to  blue.  Dyes, 
like  litmus,  whose  colors  are  affected  by  acids  and  alkalies,  are  called 
indicators.  A  substance  which  affects  an  indicator  like  an  acid  is 
said  to  have  an  acid  reaction.  A  substance  which  has  the  opposite 
effect  on  an  indicator,  behaving  like  an  alkali,  is  said  to  have  an 
alkaline  reaction.  Substances  which  do  not  affect  the  color  of 
indicators  are  said  to  be  neutral  in  reaction. 

Salts  formed  by  the  complete  neutralization  of  an  acid  by  a  base, 
in  which  all  of  the  hydrogen  of  the  acid  is  replaced  by  a  metal,  are 
called  normal  salts,  as  Na2SO4  from  H2SO4.  Such  salts  are  generally 
neutral  in  reaction.  A  salt  in  which  the  base  has  not  entirely  neu- 
tralized an  acid,  and  the  metal  has  replaced  only  a  part  of  the 
hydrogen  of  the  acid,  is  called  an  acid  salt,  as  NaHSO4.  They  are 
generally  acid  in  reaction.  A  salt  in  which  the  acid  has  not  entirely 
neutralized  the  base,  and  which  retains  a  part  of  the  oxygen,  or 
oxygen  and  hydrogen  of  the  base,  is  called  a  basic  salt,  as  BiOCl. 
A  double  salt  is  a  salt  that  contains  more  than  one  metal,  as  KNaSO4. 

Compound  Radicals. — There  are  certain  groups  of  atoms,  called 
compound  radicals,  which  behave  like  single  atoms.  They  do  not 
exist  separately,  but  one  of  them  may  be  found  in  several  different 
compounds,  and  will  pass  from  one  compound  to  another  without 
its  parts  becoming  separated.  A  compound  radical  has  a  definite 
valence,  which  is  the  difference  between  the  valences  of  its  constit- 
uent atoms.  Some  are  electro-positive  and  behave  like  metals, 
and  some  are  electro-negative  and  behave  like  non-metals.  Many 
of  these  groups  have  commonly  used  names,  like  hydroxyl  (OH)', 
ammonium  (NH4)',  and  cyanogen  (CN)'.  From  acids  there  is  an 
important  class  of  such  groups,  called  acid  radicals,  and  found  in  the 


NOMENCLATURE  17 

salts  of  the  acids,  each  consisting  of  all  of  an  acid  excepting  the 
replaceable  hydrogen.  For  example,  (NO3)',  from  HNO3,  is  found 
in  salts  of  nitric  acid;  (SO4)",  from  H2SO4,  in  salts  of  sulphuric  acid; 
(PO4)'",  from  H3PO4,  in  salts  of  phosphoric  acid,  and  so  on. 

Nomenclature. — Systematic  rules  have  been  gradually  adopted 
for  naming  chemical  compounds,  and  a  knowledge  of  them  is  impor- 
tant, as  it  enables  one  to  name  a  compound  from  its  formula,  or  to 
give  the  formula  from  its  name.  These  names  are  confusing,  as 
many  compounds,  like  water,  salt  and  marble,  have  long  been  known 
as  separate  substances  and  bear  common  names,  called  synonyms, 
which  have  no  relationship  to  their  chemical  composition;  also,  the 
rules  have  been  changed  and  improved  from  time  to  time,  but  com- 
pounds are  often  called  by  names  given  to  them  under  old  rules. 

The  only  rule  governing  the  naming  of  elements  is  that  "-um" 
indicates  a  metal,  though  there  are  several  important  metals,  like 
iron,  copper  and  gold,  whose  commonly  used  names  do  not  conform 
to  this  rule.  Some  of  the  elements,  like  gold,  sulphur  and  iron,  have 
been  known  for  a  long  time  and  their  names  have  long  been  used. 
Those  discovered  more  recently  have  been  named  from  some  striking 
property,  or  from  the  localities  in  which  they  were  first  found. 

Many  of  the  metals  vary  in  valence  and  form  two  distinct  series  of 
compounds.  These  are  distinguished  from  each  other  by  changing 
the  end  of  the  name  of  the  metal  to  -ous  for  the  compounds  in  which 
the  metal  has  the  lower  valence,  and  to  -ic  for  the  compounds  in 
which  it  has  the  higher  valence.  Examples:  compounds  of  mercury 
in  which  the  mercury  has  a  valence  of  one  are  the  mercurous  com- 
pounds, and  compounds  in  which  mercury  has  a  valence  of  two  are 
the  mercuric  compounds;  compounds  of  iron  (ferrum)  in  which  the 
iron  has  a  valence  of  two  are  the  ferrous  compounds,  and  compounds 
in  which  iron  has  a  valence  of  three  are  the  ferric  compounds. 

A  binary  compound  is  a  compound  containing  two  elements  only, 
like  HC1,  H2O,  etc.  A  binary  compound  is  called  by  the  name  of 
the  positive  element  in  the  compound  followed  by  the  name  of  the 
negative  element  with  its  last  part  changed  to  -4de.  Example: 
NaCl  is  sodium  chloride.  If  two  elements  form  two  or  more  com- 
pounds with  each  other,  these  compounds  may  sometimes  be  dis- 
tinguished by  the  rule  for  distinguishing  different  series  of  com- 
pounds of  metals.  Examples:  HgCl  is  mercurous  chloride  and 
HgCl2  is  mercuric  chloride.  Or  prefixes  are  used  to  indicate  the 
number  of  atoms  of  the  negative  element  in  a  molecule  of  the 
compound,  the  number  of  atoms  of  the  positive  element  remain- 
ing the  same.  Examples:  H2O  is  hydrogen  monoxide;  H2O2  is 
hydrogen  dioxide.  Compound  radicals  are  often  considered  as 
elements,  and  their  compounds  may  be  named  like  binary  compounds. 
Examples:  NH4C1  is  ammonium  chloride;  Hg(CN)2  is  mercuric 
2 


18  ELEMENTARY  THEORY  OF  CHEMISTRY 

cyanide.    Bases  are  named  by  the  binary  rule,  thus  CaO  is  calcium 
oxide  and  Ca(OH)2  is  calcium  hydroxide. 

Hydracids  are  named  by  taking  the  name  of  the  characteristic 
element,  changing  its  last  part  to  -ic  and  prefixing  hydro.  Examples : 
HC1  is  hydrochloric  acid;  H2S  is  hydrosulphuric  acid.  By  the 
binary  rule  these  would  be  hydrogen  chloride  and  hydrogen  sulphide, 
respectively,  and  such  names  are  sometimes  used. 

Each  negative  element  generally  forms  several  oxyacids.  In 
naming  the  oxyacids  formed  by  an  element,  one  containing  a  rela- 
tively large  number  of  oxygen  atoms  in  its  molecules  is  called  by  the 
name  of  the  element  with  its  last  part  changed  to  -ic.  Example: 
HC1O3  is  chloric  acid.  The  acid  containing  the  next  smaller  number 
of  oxygen  atoms  in  its  molecules  is  named  by  taking  the  name  of  the 
characteristic  element  and  changing  its  last  part  to  -ous.  Example: 
HC1O2  is  chlorous  acid.  If  there  is  an  acid  with  still  less  oxygen 
than  in  the  "ous"  acid  it  is  called  by  the  name  of  the  characteristic 
element,  changing  its  last  part  to  -ous  and  prefixing  hypo.  Example : 
HC1O  is  hypochlorous  acid.  If  there  is  an  acid  with  more  oxygen 
in  the  molecules  than  is  contained  in  the  "ic"  acid,  it  is  called  by  the 
name  of  the  characteristic  element,  changing  its  last  part  to  -ic 
and  prefixing  per-.  Example:  HC1O4  is  perchloric  acid. 

Salts  of  hydro-acids  are  binary  compounds  and  are  named  by  the 
binary  rule.  Example:  CaCl2  is  the  calcium  salt  of  hydrochloric 
acid  and  is  named  calcium  chloride.  Salts,  of  oxyacids  are  named 
from  the  names  of  the  acids.  A  salt  of  an  acid  whose  name  ends  in 
-ic  is  called  by  the  name  of  the  positive  element  or  radical  followed 
by  the  name  of  the  acid  with  its  last  part  changed  to  -ate.  Examples : 
KC1O3  is  the  potassium  salt  of  chloric  acid  and  its  name  is  potassium 
chlorate;  similarly  KC1O4  is  potassium  perchlorate.  A  salt  of  an 
acid  whose  name  ends  in  -ous  is  called  by  the  name  of  the  positive 
element,  followed  by  the  name  of  the  acid  with  its  last  part  changed 
to  -lie.  Examples:  KC1O2  is  the  potassium  salt  of  chlorous  acid 
and  its  name  is  potassium  chlorite;  similarly  KC1O  is  potassium 
hypochlorite.  Salts  were  formerly  named  in  reverse  order  and  the 
name  of  the  metal  modified  in  ways  no  longer  in  good  usage,  but 
these  older  names  are  often  seen.  Examples:  sodium  chloride  was 
formerly  called  chloride  of  soda;  potassium  chlorate  was  formerly 
called  chlorate  of  potash,  and  so  on. 

Acid  salts  are  generally  designated  by  the  prefix  bi-,  as  in  sodium 
bi-carbonate,  which  is  NaHCO3,  sometimes  called  sodium  hydrogen 
carbonate  or  sodium  acid  carbonate.  Basic  salts  are  commonly 
designated  by  the  prefixes  sub-  or  oxy-,  as  in  bismuth  subchloride, 
which  is  BiOCl,  sometimes  called  bismuth  oxy chloride  or  bismuth 
basic  chloride.  The  name  of  a  double  salt  contains  the  names  of 
both  metals  in  the  salt,  as  in  potassium-sodium  sulphate,  KNaSO4. 


CALIFORNIA    COLLEGF 
erf  pm  19 


There  are  some  compounds  whose  names  are  not  included  in  the 
rules  above,  but  they  may  be  learned  separately  as  the  compounds 
themselves  are  studied. 

Solution.  —  All  substances,  when  brought  into  contact  with  liquids, 
are  affected  by  the  liquid  to  a  greater  or  less  extent.  A  solid  sub- 
stance, on  being  mixed  with  a  liquid,  may  lose  its  solid  condition 
entirely  and  become  a  part  of  the  liquid,  when  it  is  said  to  be  dis- 
solved, a  solution  being  formed;  or  a  great  part  of  the  solid  may  be 
merely  suspended  in  the  liquid  in  the  form  of  small  solid  particles. 
Solids  which  dissolve  in  a  liquid  are  said  to  be  soluble  in  the  liquid, 
which  is  called  a  solvent.  Solids  which  will  not  dissolve  in  a  liquid 
are  said  to  be  insoluble  in  that  liquid,  though  this  is  a  comparative 
matter  only,  as  all  solids  will  dissolve  in  liquids  to  some  extent. 
So-called  insoluble  substances  are  only  very  slightly  soluble.  When 
a  liquid  has  dissolved  all  of  a  substance  that  it  can  dissolve  at  the 
ordinary  temperature  the  solution  is  said  to  be  saturated.  If  the 
solvent  is  heated  it  will  generally  dissolve  more  of  a  solid  than  at 
lower  temperatures,  and  if  an  excessive  amount  is  thus  dissolved  the 
solution  is  said  to  be  super-saturated,  the  excess  being  deposited  from 
the  solution  on  cooling.  All  liquids  and  gases  dissolve  in  liquids  to  a 
greater  or  less  extent,  in  the  same  way  that  solids  do.  When  two 
liquids  dissolve  each  other  in  all  proportions  they  are  said  to  be 
miscible.  Solution  is  a  phenomenon  which  is  of  great  importance 
in  natural  processes  and  in  manufacturing  operations. 

Insoluble  matter  in  a  liquid  is  called  a  sediment;  when  an  insoluble 
substance  is  formed  in  a  solution  by  chemical  reaction  or  other  means 
it  is  called  a  precipitate,  and  the  process  is  called  precipitation.  If 
the  constituents  of  a  solution  can  react  to  form  an  insoluble  sub- 
stance or  precipitate  they  will  generally  do  so.  This  is  called  the 
laiv  of  precipitation,  and  much  use  is  made  of  it  in  chemical  processes. 

To  remove  a  sediment  or  precipitate  from  a  liquid  we  generally 
employ  paper  filtration  in  analytical  work.  This  consists  of  pouring 
the  liquid  containing  the  precipitate  on  a  smooth  folded  paper 
filter  contained  in  a  glass  funnel.  The  clear  liquid  which  passes 
through  the  paper  is  called  a  filtrate.  The  folded  paper  should  not 
extend  quite  to  the  edge  of  the  funnel,  and  it  is  generally  better  to 
wet  the  paper  with  water  before  filtering  if  a  watery  liquid  is  to  be 
filtered.  If  the  filtrate  does  not  come  clear  it  will  often  do  so  if  it  is 
passed  several  times  through  the  same  filter  containing  the  precipi- 
tate. 

The  nature  of  solution  was  long  unknown,  but  it  has  been  carefully 
studied  during  recent  years.  At  first  sight  it  appears,  merely,  that 
when  a  substance  goes  into  solution  the  influences  of  the  molecules 
on  each  other  are  overcome  as  the  attraction  between  the  molecules 
in  a  solid,  or  the  repulsion  between  the  molecules  of  a  gas,  disappear. 


20  ELEMENTARY  THEORY  OF  CHEMISTRY 

Some  solutions  appear  to  be  of  this  character,  but  in  others  there  is  a 
much  more  radical  change. 

If  two  fluids,  liquids  or  gases,  which  are  miscible,  are  brought  into 
contact  with  each  other,  they  immediately  begin  to  mix,  and  this 
spontaneous  mixing  will  continue  until  a  uniform  mixture  results. 
It  is  rapid  for  gases  and  slow  for  liquids.  When  the  liquids  are  only 
superimposed,  one  on  the  other,  the  process  is  called  diffusion; 
when  they  are  separated  by  a  porous  membrane  the  process  is  called 
osmosis.  Osmosis  is  a  selective  process  when  applied  to  solutions, 
some  substances  in  solution  passing  through  the  porous  membrane 
with  the  solvent,  while  others  do  not  pass  through.  The  substances 
which  will  osmose  with  their  solvents  are  called  crystalloids,  and 
those  that  will  not  osmose  are  called  colloids.  Dialysis  is  the 
process  of  separating  crystalloids  from  colloids  by  osmosis.  Col- 
loids and  crystalloids  are  in  quite  different  conditions  in  their  solu- 
tions. Gelatin,  starch,  gums,  soap  and  rubber  are  colloids  with 
most  solvents,  also  certain  hydroxides  and  other  compounds  of 
metals,  such  as  are  found  in  milk  of  magnesia  and  dialyzed  iron. 
Colloids  do  not  form  true  solutions,  as  the  molecular  influences  are 
not  overcome.  They  appear  to  swell  and  absorb  the  solvent  to 
form  a  pasty  or  liquid  mixture.  The  boiling-  and  freezing-points 
of  the  solvent  are  not  changed,  as  they  are  in  true  solutions. 

When  crystalloids  dissolve  there  may  be  only  a  molecular  disper- 
sion, as  in  the  solution  of  sugar  in  water,  but  this  is  exceptional. 
In  nearly  all  solutions  of  crystalloids  the  molecules  of  the  dissolved 
substances  are  not  only  dispersed,  but  they  are  partially  broken 
down,  or  dissociated,  into  positive  and  negative  parts  called  ions. 
Example:  in  the  solution  of  sodium  nitrate  some  of  the  molecules 
of  NaNO3  are  separated  into  the  electro-positive  ion  Na  and  the 
electro-negative  ion  NO3.  The  relative  number  of  molecules  thus 
disassociated  varies  with  the  identity  of  the  compound,  the  solvent, 
the  temperature  and  the  strength  of  the  solution.  In  weak  solu- 
tions the  dissociation  is  relatively  greater  than  in  strong  solutions. 
The  electrical  conductivity  of  solutions  is  due  to  the  positive  and 
negative  characters  of  the  ions.  When  a  compound  is  decomposed 
by  an  electric  current  the  process  is  called  electrolysis. 

Water  of  Crystallization. — Many  chemical  compounds  tend  to  assume 
characteristic  geometric  forms,  called  crystals,  when  the  compounds 
are  deposited  from  a  solution  as  the  solvent  evaporates,  or  when 
the  compound  solidifies  after  fusion,  or  is  condensed  from  the 
gaseous  condition.  A  crystalline  compound  is  one  which  is  in  the 
form  'of  crystals,  and  an  amorphous  compound  is  one  which  is  not 
in  the  form  of  crystals.  The  crystals  of  a  compound  are  generally 
alike  in  form,  though  they  may  vary  widely  in  size,  and  they  may 
retain  some  loosely  combined  water,  which  is  called  water  of  crystal- 


CHEMICAL  ANALYSIS  21 

lization,  and  without  which  the  compound  will  not  form  crystals  or 
crystallize.  Example:  copper  sulphate  in  the  crystalline  form  has  a 
deep  blue  color  and  has  the  formula  CuSO45H2O.  If  this  is  heated 
moderately  the  water  is  driven  off  and  the  residue  of  CuSO4  is  a 
white  amorphous  powder.  Not  all  crystals,  however,  contain  water 
of  crystallization.  Example:  common  salt,  sodium  chloride,  readily 
crystallizes  and  has  the  formula  NaCl,  without  any  water  in  the 
molecules. 

The  water  of  crystallization  is  combined  in  the  molecules  of  a 
substance  and  does  not  appear  as  moisture.  If  the  water  of  crystal- 
lization is  expelled  from  a  compound  by  heating,  the  process  is 
called  exsiccation  and  the  compound  is  then  said  to  be  anhydrous. 
A  compound  which  yields  all  or  part  of  its  water  of  crystallization 
to  the  air  is  said  to  be  efflorescent.  A  substance  which  absorbs 
moisture  from  the  air  is  hygroscopic,  and  if  the  water  appears  as 
moisture,  dissolving  the  substance  partially  or  completely,  the 
substance  is  deliquescent.  Examples:  crystalline  washing  soda, 
sodium  carbonate,  has  the  formula  Na2CO310H2O  and  'effloresces 
in  dry  air,  falling  to  an  amorphous  powder.  Anhydrous  copper 
sulphate  will  absorb  water  from  moist  air,  changing  from  white  to 
blue  in  color.  Fused  calcium  chloride,  CaCl2,  is  a  deliquescent 
substance,  first  absorbing  water  to  form  CaCl26H2O,  and  then  absorb- 
ing more  water,  becoming  moist  and  ultimately  dissolving  in  the 
water. 

When  a  compound  containing  water  of  crystallization  is  dissolved 
in  water  the  water  of  crystallization  becomes  a  part  of  the  solvent 
and  is  generally  omitted  in  giving  the  formula  of  the  compound, 
but  when  a  crystalline  compound  is  not  dissolved  the  water  of 
crystallization  often  must  be  considered.  Example:  alum  crystals 
have  the  formula  KA1(S04)2  12H2O,  being  nearly  one-half  water, 
but  we  use  the  formula  KA1(S04)2  when  referring  to  alum,  unless 
there  is  a  special  reason  for  including  the  water  of  crystallization. 

Chemical  analysis  is  the  separation  of  a  substance  into  its  constit- 
uents by  chemical  processes.  Qualitative  analysis  is  the  identifica- 
tion of  the  constituents  of  a  substance.  Quantitive  analysis  is  the 
determination  of  the  relative  amounts  of  the  constituents  of  a 
substance.  Qualitative  analysis  may  be  used  to  ascertain  the 
composition  of  an  unknown  substance,  or  to  find  the  presence  or 
absence  of  impurities  in  a  known  substance. 

Each  element  and  compound  radical  has  certain  characteristic 
reactions  which  are  used  to  separate  and  identify  it.  Any  substance 
used  to  bring  about  a  reaction  is  called  a  reagent.  If  it  is  used  in 
the  form  of  a  solution  the  solution  is  called  a  test-solution.  A  list  of 
the  most  often  used  test-solutions  and  their  strengths  will  be  found 
at  the  end  of  this  book. 


THE    METALS. 

It  is  customary  to  separate  and  identify  the  metals  in  a  substance 
before  looking  for  the  acid  radicals  and  non-metals.  The  scheme 
generally  used  for  the  separation  and  identification  of  the  metals 
depends  upon  the  fact  that  a  large  majority  of  the  metals,  in  solu- 
tions of  their  compounds,  will  form  precipitates  with  hydrogen 
sulphide,  H2S,  under  various  conditions,  these  precipitates  varying 
widely  in  their  properties.  In  carrying  out  this  scheme  it  is  neces- 
sary that  the  substances  be  in  solution,  and,  at  the  beginning  of  the 
study  of  qualitative  analysis,  the  work  is  done  upon  solutions  of 
known  and  unknown  composition.  -  Afterward,  some  practice  on 
the  bringing  of  substances  into  solution  is  necessary. 

In  this  manual  the  principal  metals  are  classified  into  seven  groups, 
as  follows: 

Group  1. — Metals  precipitated  as  chlorides  by  hydrochloric  acid, 
HC1:  silver,  Ag;  mercurous  mercury,  Hg';  and  all  but  a  very  small 
amount  of  lead,  Pb. 

Group  2. — Metals  precipitated  as  sulphides  by  hydrogen  sulphide, 
H2S,  in  the  presence  of  hydrochloric  acid,  HC1,  and  whose  sulphides 
are  insoluble  in  ammonium  sulphide,  (NH4)2SX:  mercuric  mercury, 
Hg";  bismuth,  Bi;  copper,  Cu;  cadmium,  Cd;  and  the  small  amount 
of  lead  not  removed  with  group  1 . 

Group  3. — Metals  precipitated  as  sulphides  by  hydrogen  sulphide, 
H2S,  in  the  presence  of  hydrochloric  acid,  HC1,  and  whose  sulphides 
are  soluble  in  ammonium  sulphide,  (NH4)2  Sx:  arsenic,  As;  antimony, 
Sb;  tin,  Sn;  gold,  Au;  and  platinum,  Pt. 

Group  4. — Metals  precipitated  as  hydroxides  by  ammonium 
hydroxide,  NH4OH,  in  the  presence  of  ammonium  chloride,  NH4C1: 
iron,  Fe;  chromium,  Cr;  and  aluminum,  Al. 

Group  5. — Metals  precipitated  as  sulphides  by  ammonium  sul- 
phide, (NH4)2S,  or  hydrogen  sulphide,  H2S,  in  the  presence  of  am- 
monium hydroxide,  NH4OH:  cobalt,  Co;  nickel,  Ni;  manganese, 
Mn;  and  zinc,  Zn. 

Group  6. — Metals  precipitated  as  carbonates  by  ammonium 
carbonate,  (NH4)2CO3,  in  the  presence  of  ammonium  chloride, 
NH4C1,  and  ammonium  hydroxide,  NH4OH:  barium,  Ba;  strontium, 
Sr;  and  calcium,  Ca. 

Group  7. — Metals  not  precipitated  by  any  group  reagent:  mag- 
nesium, Mg;  potassium,  K;  sodium,  Na;  lithium,  Li;  and  the  com- 
pound radical  ammonium,  (NH4)'. 


LEAD  23 

The  general  reagents  used  to  precipitate  the  groups  are  called 
group  reagents  and  they  will  usually  precipitate  the  metals  of  the 
preceding  groups,  so,  in  separating  the  metals,  each  of  the  group 
reagents  must  be  used  in  exactly  the  order  given.  If  there  is  no 
precipitate  formed  when  a  group  reagent  is  added,  that  group  is 
absent  and  we  pass  on  to  the  next  group.  If  there  is  a  precipitate 
formed  the  reagent  is  slowly  added  until  precipitation  appears  to  be 
complete.  The  precipitate  is  then  filtered  out  of  the  liquid,  and 
the  filtrate  tested  first  with  a  small  amount  of  the  last  used  reagent, 
to  be  sure  that  precipitation  by  it  is  complete;  and  it  is  then  tested 
for  the  next  group. 

Nearly  all  of  the  tests  described  in  this  text-book  are  to  be  carried 
out  with  solutions.  Unless  other  directions  are  given,  from  one- 
quarter  to  one-half  an  inch  of  the  solution  .to  be  tested  is  poured  into 
a  test-tube,  for  each  test,  and  a  few  drops  of  the  reagent  are  added, 
which  are  generally  sufficient,  unless  it  is  necessary  to  completely 
remove  the  radical  being  precipitated.  Diluted  acids  and  solutions 
of  salts  are  used  as  reagents  unless  otherwise  specified. 

Success  in  the  practice  of  qualitative  analysis  depends  very  largely 
upon  the  ability  to  recognize  the  reactions  obtained.  This  ability 
is  acquired  only  by  experience,  and  much  practice  is  necessary 
upon  solutions  whose  composition  is  known,  called  known  solutions. 
After  becoming  familiar  with  the  reactions  by  work  on  known  solu- 
tions, the  student  can  undertake  the  analysis  of  unknown  solutions 
of  gradually  increasing  difficulty. 

GROUP  1. 
Metals  precipitated  as  chlorides  by  hydrochloric  acid,  HC1. 

Lead,  Pb;  mercurous  mercury,  Hg1;  silver,  Ag. 
LEAD  (PLUMBUM),  Pb11  =  207.10. 

Lead  is  largely  used  in  the  metallic  state,  alone  and  in  many 
alloys,  like  solder,  type-metal  and  pewter.  The  compounds  of  lead 
vary  in  color  some  being  white,  some  yellow,  red  or  brown.  Lead 
acetate  and  lead  nitrate  are  readily  soluble  in  water,  but  all  other 
important  salts  of  lead  are  insoluble  or  only  very  slightly  soluble 
in  water.  Lead  compounds  are  dangerous  poisons,  in  single  large 
doses  or  in  small  doses  taken  over  a  long  period  of  time. 

IMPORTANT  COMPOUNDS  OF  LEAD. 

Lead  acetate,  Plumbi  Acetas,  U.  S.  P.,  "Sugar  of  lead,"  Pb(C2H3O2)2. 
Lead  carbonate,  basic  lead  carbonate,   " white  lead,"   "cerussa," 

(PbC03)2Pb(OH)2. 
Lead  chromate,  "  chrome  yellow,"  PbCrO4. 


24  THE  METALS 

Lead  iodide,  PbI2. 

Lead  nitrate,  Pb(NO3)2. 

Lead  oxide,  Plumbi  oxidum,  U.  8.  P.,  "litharge,"  PbO. 

Red  lead  oxide,  "minium,"  "red  lead,"  Pb3O4. 

Lead  dioxide,  lead  peroxide,  PbO2. 

Lead  sulphide,  "galena,"  PbS. 

TESTS  FOR  LEAD. 

Use  a  separate  portion  of  a  solution  of  lead  acetate,  Pb(C2H302)2, 
or  of  lead  nitrate,  Pb(NO3)2,  for  each  of  the  following  tests: 

1.  Add  HC1,   obtaining  a  white  precipitate  of  lead  chloride, 
PbCl2. 

This  precipitate  is  slightly  soluble  in  water,  so  the  lead  is  not  all 
precipitated,  and  a  weak  solution  of  a  lead  salt  will  not  give  the 
reaction.  Pour  a  portion  of  the  liquid  containing  the  precipitate 
of  PbCl2  on  a  filter,  pass  hot  water  through  the  filter  and  the  precipi- 
tate will  dissolve. 

Soluble  chlorides,  as  KC1,  will  give  the  same  precipitate. 

2.  Add  KI,  obtaining  a  -yellow  precipitate  of  lead  iodide,  PbI2. 

3.  Pass  H2S  gas  through  the  lead  solution,  obtaining  a  black 
precipitate  of  lead  sulphide,  PbS. 

(NH4)2S  will  give  the  same  precipitate. 

4.  Add  H2SO4,  obtaining  a  white  precipitate  of  lead  sulphate, 
PbS04. 

Soluble  sulphates,  as  K2SO4,  will  give  the  same  precipitate. 

5.  Add  a  few  drops  of  NaOH,  obtaining  a  white  precipitate  of 
lead  hydroxide,  Pb(OH)2,  which  forms  sodium  plumbite,  Na2PbO2, 
and  dissolves  on  adding  an  excess  of  the  NaOH. 

6.  Add  K2Cr2C>7,  obtaining  a  yellow  precipitate  of  lead  chromate 
PbCr04. 

Soluble  chromates,  as  K2CrO4,  give  the  same  precipitate. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  Pb(C2H302)2 .+  HC1  = 
Pb(NO3)2  +  KC1  = 
Pb(C2H302)2  +  CaCl2  = 

2.  Pb(C2H302)2  +  KI  = 
Pb(N03)2  +  NH4I  = 

3.  Pb(NO3)2  +  H2S  = 
Pb(C2H302)2  +  (NH4)2S  = 

4.  Pb(N03)2  +  H2S04  = 
Pb(C2H302)2  +  (NH4)2S04  = 

5.  Pb(NO3)2  +  NaOH  = 
Pb(OH)2  +  NaOH  = 

6.  2Pb(C2H3O2)2  +  K2Cr2O7  +  H20  =  2PbCrO4  +  2KC2H3O2 

+  2HC2H302. 
Pb(N03)2  +  K2Cr04  = 


MERCUROUS  MERCURY  25 

MERCUROUS  MERCURY,  Hg1  =  200.6. 
HYDRARGYRUM,  U.  S.  P.,  "  QUICKSILVER." 

Mercury  is  a  silvery  white  liquid  metal  which  is  largely  used  in 
the  metallic  state  in  thermometers,  barometers  and  metallurgical 
processes.  When  thoroughly  mixed  with  some  substances,  as  chalk, 
syrup,  honey,  and  fats,  it  is  extinguished,  that  is,  divided  into  minute 
globules  which  remain  separated,  and  the  metal  in  this  condition 
has  a  dark  gray  color.  Mass  of  mercury,  mercurial  ointment  and 
mercury  with  chalk  all  contain  extinguished  metallic  mercury. 
Alloys  containing  mercury  are  called  amalgams. 

There  are  two  classes  of  mercury  salts,  both  of  which  are  impor- 
tant: the  mercurous  salts  considered  here,  and  the  mercuric  salts 
considered  in  group  2.  The  mercurous  compounds  vary  in  color, 
from  white  through  yellow  and  green  to  black.  They  are  nearly 
all  insoluble  in  water,  the  nitrate  being  the  only  important  salt 
which  is  readily  soluble.  Compounds  of  mercury  are  poisonous 
and  are  volatilized  by  heat,  with  or  without  decomposition. 

IMPORTANT  MERCUROUS  COMPOUNDS. 

Mercurous  chloride,  Hydrargyn  chloridum  mite,  U.S.  P.,  "calomel," 

HgCl. 
Mercurous  iodide,  Hydrargyri   iodidum  flavum,   U.  S.  P.,  "proto- 

iodide  of  mercury,"  "green  iodide  of  mercury,"  Hgl. 
Mercurous  nitrate,  HgNO3. 
Mercurous  oxide,  Hg2O. 

TESTS  FOR  MERCUROUS  MERCURY. 

Use  a  separate  portion  of  a  solution  of  mercurous  nitrate,  HgNO3, 
for  each  of  the  following  tests. 

1.  Add  HC1,  obtaining  a  white  precipitate  of  mercurous  chloride, 
HgCl. 

Collect  a  small  portion  of  the  precipitate  on  a  filter  and  pass  hot 
water  through  the  filter  and  the  precipitate  will  not  dissolve.  Pour 
NH4OH  through  the  filter  and  the  precipitate  will  blacken,  forming 
dimercurous  ammonium  chloride,  NH2Hg2CL 
'  2.  Add  KI,  obtaining  a  greenish-yellow  precipitate  of  mercurous 
iodide,  Hgl. 

3.  Pass  H2S  gas  through  the  solution,  obtaining  a  black  precipi- 
tate, consisting  of  mercuric  sulphide,  HgS,  and  metallic  mercury. 

(NH4)2S  will  give  the  same  precipitate. 

4.  Add  NaOH,  obtaining  a  black  precipitate  of  mercurous  oxide, 
Hg20. 

The  same  precipitate  is  obtained  with  lime  water,  Ca(OH)2.    A 


26  THE  METALS 

mixture  formed  from  calomel  and  lime  water  is  an  old-fashioned 
remedy,  called  "  black  wash/' 

§.  Pour  5  mils  of  the  mercury  solution  into  a  small  beaker. 
Place  a  small  piece  of  copper  foil  in  the  solution.  The  copper 
becomes  covered  with  a  gray  deposit  of  metallic  mercury,  which 
brightens  when  rubbed.  Dry  the  coated  copper  foil,  put  it  in  a  dry 
test-tube  and  heat  gently  over  the  flame  of  a  Bunsen  burner.  The 
mercury  will  volatilize  and  condense  on  the  cooler  part  of  the  test- 
tube  as  a  gray  deposit,  which  consists  of  small  globules  of  the  metal. 

Mercury  will  deposit  from  its  compounds  and  amalgamate  with 
many  metals  as  with  the  copper  in  this  test  and  care  must  be  taken 
not  to  allow  mercury  or  its  compounds  to  come  in  contact  with  gold 
rings. 

COMPLETE  AND  BALANCE  TriE  FOLLOWING  EQUATIONS: 

1.  HgNOs  +  HCl  = 

2HgCl  +  2NH4OH  =  NH2Hg2Cl  +  NH4C1 .+  2H2O 

2.  HgNO3  +  KI  = 
HgN03  +  NH4I  = 

3.  HgNO3  +  H2S  = 
HgN03  +  (NH4)2S  = 

4.  HgNO,  +  NaOH  = 
HgCl  +  Ca(OH)2  = 

5.  HgNOs  +  Cu  = 

SILVER  (ARGENTUM),  Ag1  =  107.12. 

Silver  is  largely  used  in  the  metallic  state,  being  generally  alloyed 
with  a  small  amount  of  copper  to  harden  it.  Its  compounds  are 
poisonous  and  unstable,  being  easily  reduced  to  the  oxide  or  metallic 
silver  by  reducing  agents,  including  organic  matter.  This  reduction 
is  rapid  in  actinic  light.  Most  of  the  important  silver  salts  are 
colorless,  but  the  iodide  is  yellow  and  the  oxide  is  black.  The 
nitrate  is  the  only  important  salt  which  is  readily  soluble  in  water. 

IMPORTANT  COMPOUNDS  OF  SILVER. 
Silver  chloride,  AgCl. 
Silver  cyanide,  AgCN. 
Silver  iodide,  Agl. 

Silver  nitrate,  Argenti  mtras,  U.  S.  P.,  "lunar  caustic,"  AgNO3. 
Silver  oxide,  Argenti  oxidum,  U.  S.  P.,  Ag2O. 
Silver  sulphate,  Ag2SO4. 

TESTS  FOR  SILVER. 

Use  a  separate  portion  of  a  solution  of  silver  nitrate,  AgNOs,  for 
each  of  the  following  tests: 


SILVER  27 

1.  Add  HC1,  obtaining  a  white  precipitate  of  silver  chloride, 
AgCl,  insoluble  in  acids,  but  very  soluble  in  ammonia  water,  NH4OH, 
forming  ammonio-silver  chloride,  (NH3)2AgCl. 

Collect  a  small  portion  of  the  precipitate  of  AgCl  on  a  filter. 
Pour  hot  water  though  the  filter  and  the  precipitate  will  not 
dissolve.  Pour  NH4OH  through  the  filter  and  the  precipitate 
will  dissolve.  Add  HNO3  to  the  ammonio-silver  solution  and  the 
precipitate,  AgCl,  will  reappear. 

Soluble  chlorides,  as  NaCl,  will  precipitate  silver  as  silver  chloride. 

2.  Add  KI,  obtaining  a  light  yellow  precipitate  of  silver  iodide, 
Agl,  slowly  soluble  in  an  excess  of  the  reagent  as  potassium  silver 
iodide,  KAgI2. 

3.  Add  KCN,  obtaining  a  white  precipitate  of  silver  cyanide, 
AgCN,  readily  soluble  in  an  excess  of  the  reagent  as  potassium  silver 
cyanide,  KAg(CN)2. 

4.  Pass  H2S  gas  through  the  silver  solution,  obtaining  a  black 
precipitate  of  silver  sulphide,  Ag2S. 

(NH4)2S  will  give  the  same  precipitate. 

5.  Add  NaOH,  obtaining  a  brown  precipitate  of  silver  oxide, 
Ag2O,  insoluble  in  an  excess  of  the  reagent,  but  readily  soluble  in 
nitric  and  acetic  acids  and  ammonia  water. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  AgN03  +  HCL  = 
Ag2SO4  +  NaCl  = 

AgCl  +  NH4OH  =  (NH3)2AgCl  + 
(NH3)2AgCl  +  2HNO3  =  AgCl  +  2NH4N03 

2.  AgNOs  +  KI  = 
AgC2H302  +  NH4I  = 
Agl  +  KI  = 

3.  AgNOs  +  KCN  = 
Ag2S04  +  KCN  = 

4.  AgNOs  +  H2S  = 
Ag2S04  +  (NH4)2S  = 

5.  AgN03  +  NaOH  = 
Ag2O  +  HNO3  = 

SAMPLE  NUMBER  1. 
Analysis  of  a  solution  for  group  1  of  the  metals. 

The  solution  may  contain  salts  of  lead,  mercurous  mercury,  or 
silver,  or  any  two  of  these  metals,  or  all  three  of  them. 

Dip  a  small  piece  of  red  litmus  paper  in  the  solution.  If  the  color 
of  the  litmus  is  not  affected  the  solution  is  neutral  or  acid  in  reaction, 


28  THE  METALS 

as  it  should  be.  If  the  color  changes  to  blue  add  HNOs  T.S.  slowly, 
with  stirring,  until  the  liquid  will  change  the  color  of  blue  litmus 
paper  to  red. 

A.  To  about  10  mils  of  the  solution  contained  in  a  test-tube,  add 
a  few  drops  of  HC1  T.S.    If  no  precipitate  forms,  silver  and  mercur- 
ous  mercury  are  absent,  and  lead,  if  present,  is  in  very  small  amount, 
which  will  be  detected  in  group  2.     If  a  white  precipitate  forms  it 
consists  of  chlorides  of  metals  of  group  1.     Continue  to  add  the 
HC1  T.S.  until  the  precipitation  appears  to  be  complete.    Filter 
the  liquid  and  test  the  clear  filtrate  by  adding  a  few  drops  of  HC1 
T.S.    If  this  produces  a  precipitate  add  more  HC1  T.S.,  filter  through 
the  same  filter  paper  and  test  the  last  portion  of  the  filtrate  in  the 
same  manner  as  before,  and  repeat  the  filtration  and  testing  of  the 
filtrate  until  the  HC1  T.S.  produces  no  precipitate  in  the  filtrate, 
when  precipitation  is  complete.    A  small  amount  of  lead,  however, 
remains  in  the  solution,  if  this  metal  is  present. 

B.  Heat  some  water  to  boiling  in  a  beaker  supported  on  wire 
gauze  over  a  gas  flame.    Pour  20  mils  of  this  hot  water  through  the 
precipitate  on  the  filter  paper.    If  lead  is  present,  PbCl2  dissolves. 
Test  separate  portions  of  the  filtrate  for  lead  with  KI  T.S.,  H2S, 
H2SO4  T.S.,  and  K2Cr2O7  T.S.  as  described  on  page  17,  remembering 
that  the  precipitates,  if  any,  are  small  in  amount  because  the  solu- 
tion of  lead  is  very  weak. 

If  lead  is  found  pass  50  mils  or  more  of  boiling  water  through  the 
precipitate  on  the  filter  paper  to  wash  out  the  remainder  of  the 
PbCl2.  When  the  PbCl2  is  all  washed  out  a  portion  of  the  filtrate 
will  not  give  a  white  precipitate  with  H2SO4  T.S.,  with  which  it 
should  be  tested. 

If  the  precipitate  all  dissolves  in  hot  water  it  consisted  of  PbCl2 
alone,  and  silver  and  mercurous  mercury  are  absent. 

C.  If  a  precipitate  remains  on  the  filter  paper  it  consists  of  AgCl 
or  HgCl,  or  both.    Pour  a  small  amount  of  NH4OH  T.S.  through  the 
precipitate  on  the  filter.    If  the  precipitate  blackens,  the  presence 
of  mercurous  mercury  is  shown,  as  described  on  page  25. 

D.  If  silver  is  present  the  AgCl  dissolves  in  the  NH4OH  T.S. 
Add  HNOs  T.S.  to  the  filtrate  until  it  is  acid  in  reaction  to  litmus. 
A  white  precipitate  shows  the  presence  of  silver,  as  described  on 
page  27. 

GROUP  2. 

Metals  precipitated  by  hydrogen  sulphide,  H2S,  in  presence  of 
hydrochloric  acid,  HC1,  and  whose  sulphides  are  insoluble  in 
ammonium  sulphide,  (NH4)2  Sx. 

Mercuric  mercury,  Hg";  (lead,  Pb);  bismuth,  Bi;  copper,  Cu; 
cadmium,  Cd. 


MERCURIC  MERCURY  29 


MERCURIC  MERCURY   (HYDRARGYRUM),   Hgu  =  200.6. 

Oxidizing  agents  readily  change  mercurous  compounds  to  mer- 
curic compounds  and  reducing  agents  readily  change  mercuric  com- 
pounds to  mercurous  compounds.  As  mercuric  compounds  are  much 
more  poisonous  than  mercurous  compounds,  care  must  be  taken  not 
to  prescribe  or  dispense  any  mercurous  compound  for  medicinal  use 
in  such  a  way  that  it  might  be  oxidized  to  the  mercuric  condition. 

Mercuric  compounds  are  variously  colored,  being  white,  yellow, 
red  or  black.  Most  of  them  are  insoluble  or  only  slightly  soluble  in 
water,  but  mercuric  chloride,  mercuric  nitrate,  and  mercuric  cyanide 
are  soluble. 

IMPORTANT  MERCURIC  COMPOUNDS. 

Mercuric-ammonium  chloride,  Hydrargyrum  ammoniatum,  U.  S.  P., 

ammoniated  mercury,  "  white  precipitate,"  NH2HgCl. 
Mercuric  bromide,  HgBr2. 
Mercuric    chloride,    Hydrargyri   chloridum  corrosivum,    U.  S.  P., 

"  corrosive  sublimate,"  HgCl2. 
Mercuric  cyanide,  Hg(CN)2. 
Mercuric  iodide,  Hydrargyri  iodidum  rubrum,  U.  S.  P.,  "biniodide 

of  mercury,"  HgI2. 
Mercuric  nitrate,  Hg(NO3)2. 

Yellow  mercuric  oxide,  Hydrargyri  oxidum  flavum,  U.  S.  P.,  HgO. 
Red  mercuric  oxide,  Hydrargyri,  oxidum  rubrum,  U.  S.  P.,  "red 

precipitate,"  HgO. 
Mercuric  sulphate,  HgSO4. 

Mercuric  subsulphate,  "turpeth  mineral,"  (HgO)2HgS04. 
Black  mercuric  sulphide,  "ethiops  mineral,"  HgS. 
Red  mercuric  sulphide,  " cinnabar,"  "vermilion,"  HgS. 

TESTS  FOR  MERCURIC  MERCURY. 

Use  a  separate  portion  of  a  solution  of  a  mercuric  salt,  as  mercuric 
chloride,  HgCl2,  for  each  of  the  following  tests : 

1.  Pass  H2S  through  the  mercuric  solution.    A  black  precipitate 
of  mercuric  sulphide,  HgS,  is  formed,  but  several  other  colors,  rang- 
ing from  white  through  yellow,  orange  and  brown,  generally  appear 
before  the  final  black  precipitate.    This  range  of  colors  is  due  to  the 
temporary  formation  of  compounds  of  HgS  with  the  original  mer- 
curic salt. 

2.  Add  NaOH  T.S.  to  the  mercuric  solution.    An  excess  of  the 
alkali  precipitates  yellow  mercuric  oxide,  HgO.    When  the  mercuric 
salt  is  in  excess,  the  precipitate  is  brown  and  consists  of  a  basic 
salt  of  mercury.    The  same  precipitate  is  obtained  with  lime  water, 
Ca(OH)2,  and  the  mixture  formed  is  the  old  fashioned  "yellow 
wash." 


30  THE  METALS 

3.  Add  NH4OH  T.S.   to   the  mercuric  solution.    A  white  pre- 
cipitate of  mercurammonium  chloride,  NH2HgCl,  is  formed.     Add 
HC1  T.S.    The  precipitate  will  dissolve. 

4.  Slowly  add  KI  T.S.  to  the  mercuric  solution.    A  precipitate 
of  mercuric  iodide,  HgI2,  first  yellow,  then  red,  is  produced,  which 
forms  potassium-mercuric  iodide,  (KI)2HgI2,  and  dissolves  in  an 
excess  of  the  reagent. 

5.  Slowly  add  SnCl2  T.S.  to  the  mercuric  solution.    At  first  the 
precipitate  formed  is  white  and  consists  of  mercurous  chloride, 
HgCl.    When  the  reagent  is  in  excess,  the  precipitate  is  gray  and 
consists  of  finely  divided  metallic  mercury. 

6.  Acidulate  a  small  amount  of  the  mercuric  solution  in  a  small 
beaker  with  HC1  T.S.  and  introduce  a  small  piece  of  copper  foil 
into  the  liquid.    Mercury  is  deposited  on  the  copper,  just  as  was 
described  for  mercurous  compounds  on  page  18. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  HgCl2  +  H2S  = 
Hg(N03)2  +  (NH4)2S  = 

2.  HgCl2  +  NaOH  = 
HgCl2  +  Ca(OH)2  - 

3.  HgCl2  +  NH4OH  =  NH2HgCl  + 
NH2HgCl  +  HC1  - 

4.  HgCl2  +  KI  = 
Hg(N03)2  +  NH4I  = 

5.  HgCl2  +  SnCl2  = 
HgCl  +  SnCl2  = 

6.  HgCl2  +  Cu  = 

BISMUTH,  BiHi  v  =  208.0. 

Bismuth  is  a  hard  and  brittle  crystalline  metal  which  is  a  con- 
stituent of  some  easily  fusible  alloys  and  is  used  in  manufacturing  its 
compounds.  The  salts  of  bismuth  are  generally  white  in  color, 
but  the  oxide  is  yellow  and  the  sulphide  is  black.  Bismuth  chloride 
and  bismuth  nitrate  are  soluble  in  water  with  some  decomposition. 
The  other  important  compounds  are  insoluble.  Bismuth  com- 
pounds in  solution  readily  form  insoluble  basic  compounds  con- 
taining the  monad  radical  bismuthyl,  (BiO)'.  Some  organic  acids 
or  large  amounts  of  mineral  acids  prevent  this  precipitation.  Bis- 
muthic  compounds  are  unimportant. 

IMPORTANT  COMPOUNDS  OF  BISMUTH. 

Bismuth  chloride,  BiCl3. 

Bismuth  subcarbonate,  U.  S.  P.,  mainly  (BiO)2CO3H2O. 

Bismuth  nitrate,  Bi(NO3)3. 

Bismuth  oxide,  Bi2O3. 

Bismuth  subnitrate,  U.  S.  P.,  mainly,  Bi(OH)2NO3. 


COPPER  31 

TESTS  FOR  BISMUTH. 

Use  a  separate  portion  of  a  solution  of  either  bismuth  nitrate, 
Bi(NO3)3,  or  bismuth  chloride,  BiCl3,  for  each  of  the  following  tests: 

1.  Pour  2  mils  of  the  bismuth  solution  into  10  mils  of  water 
contained  in  a  test-tube.    A  white  precipitate  of  bismuth  subnitrate, 
BiONO3,  or  of  bismuth  subchloride,  BiOCl,  is  produced  if  the  liquid 
is  not  too  strongly  acid. 

2.  Pass  H2S  through  the  bismuth  solution.    A  black  precipitate 
of  bismuth  sulphide,  Bi2S3,  is  produced. 

3.  Add  NaOH  T.S.  to  the  bismuth  solution.    A  white  precipi- 
tate of  bismuth  hydroxide,  Bi(OH)3,  is  formed.     Boil  the  liquid 
containing  the  precipitate.    The  precipitate  changes  to  yellow  bis- 
muth oxide,  Bi2O3.    If  reducing  agents  are  present  the  precipitate 
becomes  black  and  then  consists  of  metallic  bismuth. 

4.  Add  Na2CO3  T.S.  to  the  bisniuth  solution.    A  white  precipi- 
tate of  bismuth  subcarbonate,  (BiO)2CO3,  is  produced. 

5.  Add  K2Cr2O7  T.S.  to  the  bismuth  solution.    A  yellow  pre- 
cipitate of  bismuth  subchromate,  (BiO)2CrO4  is  formed,  which  is 
insoluble  in  NaOH  T.S.    Solutions  of  chromates,  as  K2CrO4,  give 
the  same  precipitate. 

6.  Slowly  add  NaOH  T.S.  to  1  mil  of  SnCI2  T.  S.  until  the  pre- 
cipitate first  formed  dissolves,  forming  sodium  stannite,  Na2Sn02. 
Add  a  few  drops  of  the  bismuth  solution  to  the  Na2SnO2  solution. 
A  black  precipitate  of  metallic  bisjmuth,  Bi,  is  produced. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  Bi(NO3)3  +  H,O  = 
BiCl3  +  H20  = 

2.  BiCl3  +  H2S  = 
Bi(N03)3  +  (NH4)2S  = 

3.  Bi(N03)3  +  NaOH  = 
BiCl3  +  KOH  = 
Bi(OH)3  +  heat  = 

4.  2Bi(NO3)3  +  3Na2CO3  =  (BiO)2C03  +  6NaN08  +  2  CO2 
BiCl3  +  K2C03  = 

5.  BiCl3  +  K2Cr2O7  +  5H2O  =  2(BiO)2CrO4  +  2KC1  +  10HC1 
Bi(N03)3  +  K2Cr04  +  2H2O  =  (BiO)2CrO4  +  2KNO3  +  4HN03 

6.  4NaOH  +  SnCl2  =  Na2SnO2  +  2NaCl  +  2H2O 

2BiCl3  +  3Na2Sn02  +  GNaOH  =  2Bi  +  GNaCl  +  3Na2SnO3 
+  3H2O 

COPPER  (CUPRUM),  Cu1-"  =  63.57 

Copper  is  a  red  metal  which  is  largely  used  in  the  metallic  state, 
alone  and  in  many  alloys,  of  which  brass,  composed  of  copper  and 


32  THE  METALS 

zinc,  is  the  most  important.  Many  of  the  compounds  of  copper 
are  green  or  blue  in  color  and  they  impart  these  colors  to  their 
solutions.  The  cuprous  compounds  are  unstable  and  unimportant. 
Copper  sulphate,  copper  chloride  and  copper  nitrate  are  readily 
soluble  in  water.  The  other  important  compounds  of  copper  are 
insoluble  or  only  slightly  soluble  in  water. 

IMPORTANT  COMPOUNDS  OF  COPPER. 

Basic  copper  acetate,  "  verdigris,"  Cu2O(C2H3O2)2. 
Copper  aceto-arsenite,  "  Paris  green,"  Cu(C2H3O2)2  +  3Cu(AsO2)2. 
Copper  arsenite,  "Scheele's  green,"  CuHAsO3. 
Copper  chloride,  cupric  chloride,  CuCl22H2O. 
Copper  nitrate,  cupric  nitrate,  Cu(NO3)23H20. 
Cuprous  oxide,  Cu2O. 
Cupric  oxide,  CuO. 

Copper  sulphate,  cupric  sulphate,  Cupri  sulphas,  U.  S.  P.,  "blue 
vitriol,"  CuSO45H20. 

TESTS  FOR  COPPER. 

A  separate  portion  of  a  solution  of  a  copper  salt,  as  copper  sul- 
phate, CuSO4,  should  be  used  for  each  of  the  following  tests: 

1.  Pass  H2S  through  the  copper  solution.    A  black  precipitate 
of  copper  sulphide,  CuS,  is  formed.    This  precipitate  is  soluble  in 
nitric  acid  and  in  a  solution  of  potassium  cyanide. 

2.  Add  NaOH  T.S.  to  the  copper  solution.    A  light  blue  precipi- 
tate of  copper  hydroxide,  Cu(OH)2,  is  produced,  which  is  insoluble 
in  an  excess  of  the  reagent.    Boil  the  liquid  containing  the  precipi- 
tate.   The  copper  hydroxide  is  decomposed  into  black  cupric  oxide, 
CuO,  and  water. 

If  the  liquid  contains  citrates  or  tartrates  or  certain  other  organic  substances, 
the  copper  hydroxide  will  dissolve  in  an  excess  of  the  alkali,  forming  a  deep  blue 
solution  from  which  red  cuprous  oxide,  Cu2O  will  precipitate  on  boiling  if  reducing 
agents  are  present.  This  behavior  is  the  basis  of  Fehling's  test  for  sugar  and 
other  similar  tests. 

3.  Slowly  add  NH4OH  T.S.  to  the  solution  of  a  copper  salt.    A 
light  blue  precipitate  of  copper  hydroxide,  Cu(OH)2,  is  produced 
and  immediately  dissolves  in  an  excess  of  the  reagent,  forming  a 
deep  blue  solution  containing  ammonio-copper  sulphate,  Cu(NH3)4- 
SO4,  or  a  similar  compound.    Slowly  add  KCN  T.S.  to  the  blue 
liquid.     Potassium  cuprocyanide,  K3Cu(CN)4,  is  formed  and  the 
liquid  becomes  colorless.    H2S  does  not  precipitate  the  copper  from 
this  solution,  as  copper  sulphides  are  soluble  in  a  solution  of  KCN. 

4.  Add  K4Fe(CN)6  T.S.  to  the  solution  of  a  copper  salt.    A  red- 
brown  precipitate  of  copper  ferrocyanide,  Cu2Fe(CN)6,  is  produced, 
which  is  insoluble  in  diluted  acids. 


COLLEGf 
Oi   PHARMACY 

CADMIUM  33 

5.  Dip  a  piece  of  platinum  wire  in  the  copper  solution  and  heat 
in  the  blue  flame  of  a  Bunsen  burner.  The  flame  will  be  colored 
green  or  greenish-blue. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 


1.  CuS04 
Cu(N03)2  +  H2S  = 

2.  CuSO4  +  NaOH  = 
Cu(C2H3O2)2  +  KOH  = 
Cu(OH)2  +  heat  = 

3.  CuSO4  +  NH4OH  = 

CuS04  +  4NH4OH  =  Cu(NH3)4SO4  +  H2O 
2Cu"(NH3)4S04  +  10KCN  +  8H2O 

=  2K3Cu'(CN)4  +  2K2S04  +  8NH4OH  +  (CN)2 
(CN)2  +  2NH4OH  =  NH4CNO  +  NH4CN  +  H2O 

4.  CuS04  +  K4Fe(CN)6  = 
Cu(NO3)2  +  K4Fe(CN)6  = 

CADMIUM,  Cd11  =  112.4. 

Cadmium  is  a  soft,  white,  crystalline  metal  resembling  zinc  and  tin. 
It  is  a  constituent  of  some  easily  fusible  alloys.  Cadmium  salts  are 
generally  white  in  color,  but  the  sulphide  is  yellow  and  the  oxide  is 
brown.  The  sulphate,  the  nitrate  and  the  halogen  salts  are  soluble 
in  water.  Most  of  the  other  cadmium  salts  are  insoluble  in  water. 

IMPORTANT  COMPOUNDS  OF  CADMIUM. 

Cadmium  bromide,  CdBr24H20. 
Cadmium  chloride,  CdCl22H2O. 
Cadmium  iodide,  CdI2. 
Cadmium  nitrate,  Cd(NO3)24H2O. 
Cadmium  oxide,  CdO. 
Cadmium  sulphate,  (CdSO4)38H2O. 
Cadmium  sulphide,  CdS. 

TESTS  FOR  CADMIUM. 

Use  a  separate  portion  of  a  solution  of  a  cadmium  salt,  as  CdSO4, 
for  each  of  the  following  tests  : 

1.  Pass  H2S  through  the  cadmium  solution.    A  yellow  precipi- 
tate of  cadmium  sulphide,  CdS,  will  form.    This  precipitate  is  soluble 
in  hot  diluted  H2SO4,  but  is  not  dissolved  by  other  diluted  acids,  nor 
by  solutions  of  alkalis,  sulphides  or  cyanides. 

2.  Add  NaOH  T.S.  to  the  cadmium  solution.    A  white  precipi- 
tate of  cadmium  hydroxide,  Cd(OH)2,  will  form,  which  is  insoluble 

3 


34  THE  METALS 

in  an  excess  of  the  reagent.    The  same  precipitate  will  form  with 
NH4OH  T.S.,  but  is  soluble  in  an  excess  of  this  reagent. 

3.  Add  Na2CO3  T.S.  to  the  cadmium  solution.    A  white  precipi- 
tate of  cadmium  carbonate,  CdCO3,  is  formed,  which  is  not  soluble 
in  an  excess  of  the  reagent,  but  which  will  dissolve  in  diluted  acids. 

4.  Slowly  add  KCN  T.S.  to  the  cadmium  solution.    A  white  pre- 
cipitate of  cadmium  cyanide,  Cd(CN)2,  will  be  produced  and  this  will 
form  potassium  cadmium  cyanide,  K2Cd(CN)4,  and  dissolve  in  an 
excess  of  the  reagent.     Pass  H2S  through  the  solution.     Yellow 
cadmium  sulphide,  CdS,  will  be  precipitated. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.    CdS04 


Cd(N03)2  +  H2S  = 

2.  CdS04  +  NaOH  = 
Cd(N03)2  +  NH4OH  = 

3.  Cd(NO3)2  +  Na2CO3  = 
CdS04  +  (NH4)2C03  = 

4.  CdS04  +  KCN  = 
Cd(CN)2  +  KCN  = 
K2Cd(CN)4  +  H2S  = 

SAMPLE  NUMBER  2. 

Analysis  of  a  solution  for  group  2  of  the  metals. 

The  solution  may  contain  salts  of  mercuric  mercury,  lead,  bis- 
muth, copper  or  cadmium,  or  of  several  of  these  metals. 

Add  a  few  drops  of  HC1  T.S.  to  about  10  mils  of  the  sample,  warm 
and  pass  H2S  through  the  acidulated  liquid.  If  no  precipitate  is 
produced,  the  metals  of  group  2  are  absent.  If  a  precipitate  is  pro- 
duced, continue  to  pass  H2S  through  the  liquid  until  precipitation 
is  complete. 

Collect  the  precipitate  on  a  filter  and  test  the  filtrate  by  passing 
H2S  through  it  for  a  few  minutes.  If  any  additional  precipitate  is 
formed,  continue  to  pass  H2S  through  the  liquid  for  some  time  and 
then  filter  through  the  same  filter  paper,  testing  the  new  filtrate  in 
the  same  manner  as  before.  The  metals  of  this  group  are  all  in  the 
precipitate  on  the  filter  paper  and  the  filtrate  may  be  discarded, 
unless  it  is  to  be  tested-  for  the  metals  of  subsequent  groups. 

A.  Wash  the  precipitate  on  the  filter  with  water  and  allow  the 
water  to  drain  out,  discarding  the  washings.  Pour  the  same  portion 
of  about  5  mils  of  hot  HNO3  T.S.  through  the  filter  several  times. 
Any  white  or  light  colored  residue  insoluble  in  HNO3  T.S.  may  be 
disregarded.  A  dark  colored  residue,  insoluble  in  HNO3  T.S.  should 


ANALYTICAL  TABLES  35 

be  examined  for  mercury  by  B.    The  filtrate  is  to  be  tested  for  the 
other  members  of  the  group. 

B.  If  a  dark  colored  residue  is  left  that  will  not  dissolve  in 
HNO3  T.S.,  mercury  is  indicated.    Verify  by  puncturing  the  filter 
paper  and  washing  the  residue  into  a  test-tube  with  about  3  mils  of 
nitrohydrochloric  acid.    Boil  the  liquid  until  it  ceases  to  smell  of 
chlorine,  dilute  with  about  5  mils  of  water,  filter  and  add  SnCl2  T.S. 
A  white  or  gray  precipitate  shows  the  presence  of  mercury. 

C.  Pour  the  filtrate  obtained  in  A  into  a  beaker,  add  about  2 
mils  of  concentrated  H2SO4  and  boil  until  white  fumes  are  given  off. 
Cool  and  dilute  with  about  5  mils  of  water.    A  white  precipitate 
indicates  lead.     If  a  precipitate  forms,  filter  and  save  the  filtrate  to 
be  tested  for  the  remaining  members  of  the  sub-group. 

D.  To  verify  the  presence  of  lead,  wash  the  precipitate  on  the 
filter  paper  with  water,  allow  the  water  to  drain  out  and  pass  about 
5  mils  of  NH4C2H3O2  T.S.  through  the  paper  several  times.    Add 
K2Cr2O7  T.S.  to  the  filtrate.   A  yellow  precipitate  shows  the  presence 
of  lead. 

E.  To  the  liquid  or  filtrate  from  C,  add  an  excess  of  NH4OH  T.S., 
shown  by  a  strong  odor  of  ammonia  which  persists  after  shaking  the 
liquid.    A  white  precipitate  indicates  the  presence  of  bismuth.    If  a 
precipitate  forms,  filter  the  liquid  and  preserve  the  filtrate  to  be 
tested  for  Cu  and  Cd. 

F.  To  verify  the  presence  of  bismuth,  wash  the  precipitate  on 
the  filter  paper  with  water,  allow  it  to  drain  and  moisten  it  with  a 
few  drops  of  NaOH  T.S.,  followed  by  a  few  drops  of  SnCl2  T.S.    A 
brown  or  black  coloration  shows  the  presence  of  bismuth. 

G.  If  the  liquid  or  filtrate  from  E  is  blue,  copper  is  present. 
Verify  the  presence  or  absence  of  copper  by  adding  an  excess  of 
HC2H3O2T.S.  followed  by  K*Fe(CN)6  T.S.  to  a  portion  of  the 
liquid.    A  red    coloration    or   precipitate    shows   the   presence  of 
copper. 

II .  If  the  liquid  or  filtrate  from  E  is  colorless,  pass  H2S  through 
it.  A  yellow  precipitate  shows  the  presence  of  cadmium.  If  the 
liquid  is  blue,  slowly  add  KCN  T.S.  to  a  portion,  until  the  blue 
color  disappears  and  then  pass  H2S  through  the  liquid.  A  -yellow 
precipitate  shows  the  presence  of  cadmium. 


ANALYTICAL  TABLES. 

Condensed  directions  for  carrying  out  analytical  processes  are 
frequently  given  in  tabular  form,  and  such  tables  are  very  con- 
venient. They  are  not  often  complete  enough,  however,  to  be  fol- 
lowed successfully  by  inexperienced  workers,  so  they  should  be  used 


36 


THE  METALS 


only  by  those  who  have  become  familiar  with  the  processes  by  fol- 
lowing detailed  directions. 

In  using  the  tables  in  this  book  the  following  points  should  be 
kept  in  mind,  constantly: 

Where  chemical  formulas  are  given  for  reagents,  the  test  solutions 
should  be  used,  unless  directed  otherwise. 

Care  should  be  taken  to  use  sufficient  of  each  reagent  to  accom- 
plish its  purpose,  without  using  a  large  excess,  and  to  test  all  filtrates 
after  removing  a  metal  or  group  by  precipitation. 

Only  positive  results  are  given  in  the  tables,  as  if  all  of  the  metals 
included  are  present.  When  a  negative  result  is  obtained  on  testing 
for  any  metal  or  group,  it  is  absent  and  one  should  pass  on  to  the 
next  test,  often  omitting  filtrations  or  other  processes  which  would 
be  necessary  if  positive  results  were  obtained. 

SAMPLE  NUMBER  3. 
Analysis  of  a  solution  for  groups  1  and  2  of  the  metals. 

Examine  the  solution  by  the  following  tables,  referring  when 
necessary  to  the  detailed  directions  for  each  group  given  on  preced- 
ing pages. 


The  solution  should  be  neutral  or  acid  to  litmus.    If  alkaline,  add  HNO3  to 
the  portion  to  be  examined  until  the  reaction  is  acid. 
Add  HC1  as  long  as  a  precipitate  is  produced  and  filter. 


Precipitate — Pb,  Hgi,  Ag 

Examine  for  metals  of  group  1  by  I. 


Filtrate— Hgii,  Pb,  Bi,  Cu,  Cd 

Warm  and  pass  H2S  as  long  as  a  ppt.  is 

produced.    Filter.     Examine  for 

metals  of  group  2  by  II. 


I. 

Examination  of  any  precipitate  produced  by  HC1  in  a  neutral  or 
acid  solution. 


Wash  the  precipitate  on  the  filter  with  cold  water,  discarding  the  washings. 
Pass  a  portion  of  hot  water  through  the  washed  precipitate  on  the  filter. 


Filtrate— PbCl2 

Add  K2Cr2O7  and  cool. 

Yellow  ppt.,  PbCrO4, 

shows 


lead 


Residue— HgCl,  AgCl 

Pass  the  same  portion  of  NH4OH  through  the  residue 
on  the  filter  several  times 


Residue— NH2Hg2Cl 
black,  shows 


mercurous  mercury 


Filtrate— (NH3)3(AgCl)2 

Add  an  excess  of  HNO3, 

white  ppt.,  AgCl,  shows 

silver 


GOLD 


37 


II. 

Examination  of  a  precipitate  produced  by  H2S  in  an  acid  liquid, 
after  group  1  of  the  metals  has  been  removed  from  a  solution. 

Wash  the  residue  on  the  filter  with  water  and  drain,  discarding  the  washings. 
Pour  a  portion  of  hot  HNO3  through  the  filter  several  times. 


Residue—  HgS 

Filtrate—  Pb(NO3)2,  Bi(NO3)3,  Cu(NO3)2,  Cd(NO3)2 

Black. 

Add  cone.  H2SO4  and  boil  until  white  fumes  are  given  off,  dilute 

Dissolve  in 

with  H2O,  cool  and  filter. 

nitrohydro- 
chloric     acid, 

Precipitate  — 

TpUOfl 

Filtrate—  Bi2(SO4)3,  CuSO4,  CdSO4 

boil  to  expel 

ruB\J± 

Add  an  excess  of  NH4OH  and  filter. 

Cl,  dilute  with 
water,  filter 

White. 
Dissolve  with 

Precipitate  — 

Filtrate— 

and  add  SnCl2. 

NH4C2H3O2 

Cu(NH3)4SO4,  Cd(NH3)4SO4 

A  white  or 
gray  ppt. 

and  add        !  White.     Add 
K2Cr2O7.           NaOH  and 

Divide  into  two  portions. 

HgCl+Hg, 

Yellow  ppt., 

SnCl2  on 

To  one  portion 

Pass    H2S 

shows          PbCrO4,  shows 

filter.   Brown 

add  an  excess 

through  the 

or  black  col- 

of HC2H3O2 

other  portion, 

oration,  Bi2Os, 

and  then 

first    decoloriz- 

shows 

K4Fe(CN)6. 

ing  with  KCN 

A  red  ppt.  or 

if  Cuis  present. 

coloration, 

A  yellow  ppt., 

Cu2Fe(CN)6, 

CdS,  shows 

mercuric 

shows 

mercury 

lead 

bismuth 

copper 

cadmium 

GROUP  3. 

Metals  precipitated  as  sulphides  by  hydrogen  sulphide,  H2S, 
in  the  presence  of  hydrochloric  acid,  HC1,  and  whose  sulphides  are 
soluble  in  ammonium  sulphide,  (NH4)2SX. 

Gold,  Au;  Platinum,  Pt;  Arsenic,  As;  Antimony,  Sb;  Tin,  Sn. 


GOLD  (AURUM)  Au1 


197.2. 


Gold  is  a  soft  yellow  metal  that  is  used  most  largely  in  the  metallic 
state  in  coins,  jewelry,  etc.  It  is  permanent  in  the  air  and  does  not 
dissolve  in  any  single  acid,  but  is  dissolved  by  nitro-hydrochloric 
acid  and  other  liquids  containing  free  chlorine,  bromine  or  iodine. 
The  metal  is  also  attacked  by  fused  alkalies  and  fused  potassium 
nitrate.  Gold  compounds  are  poisonous  and  unstable.  The  only 
important  compounds  of  gold  are  the  auric  salts  with  the  halogen 
elements,  which  are  soluble  in  water  and  brown  or  yellow  in  color. 
The  aurous  compounds  are  unimportant. 


38  THE  METALS 

IMPORTANT  COMPOUNDS  OF  GOLD. 

Gold  chloride,  AuCl3. 

Gold  and  sodium  chloride,  Auri  et  Sodii  Chloridum,  U.  S.P.,  AuCl3+ 

NaCl. 

Bromauric  acid,  N.  F.,  HAuBr4  +  5H2O. 
Chlorauric  acid,  HAuCU  +  4H2O. 

TESTS  FOR  GOLD. 

Use  a  separate  portion  of  a  solution  of  gold  chloride,  AuCl3,  for 
each  of  the  following  tests : 

1.  Pass  H2S  gas  through  the  gold  solution,  obtaining  a  brownish- 
black  precipitate  of  gold  sulphide,  Au2S3.    Collect  the  precipitate  on 
a  filter  and  pass  yellow  ammonium  sulphide  solution  through  the 
filter.    The  precipitate  will  dissolve. 

2.  Add  NaOH  T.S.     A  brown  precipitate  of  gold  hydroxide, 
Au(OH)3,  soluble  in  an  excess  of  the  reagent  will  be  formed. 

3.  Add  2  drops  of  nitric  acid  and  then  FeSO4  T.S.    A  brown 
precipitate  of  metallic  gold  will  be  formed. 

4.  Add  SnCl2  T.S.  and  let  stand.   A  purple  precipitate  of  "  Purple 
of  Cassius"  is  slowly  formed.    This  is  a  substance  of  indefinite  com- 
position, containing  aurous  oxide,  Au^O,  and  oxides  of  tin. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  AuCl3  +  H2S  = 

2.  AuCl3  +  NaOH  = 

3.  2AuCl3  +  6FeS04  =  2  Au  +  2FeCl3  +  2Fe2(SO4)3 

PLATINUM,  Ptu  iv  =  195.2. 

Platinum  is  a  heavy  white  metal  with  a  high  melting-point.  It  is 
permanent  in  the  air  and  is  not  attacked  by  any  single  acid.  Plati- 
num vessels  are  used  in  chemical  work  and  care  should  be  taken  not 
to  allow  them  to  come  in  contact  with  free  chlorine  or  bromine,  nor 
to  ignite  alkalies  in  them,  nor  mixtures  that  contain  free  iodine, 
sulphur,  phosphorus  or  metals,  or  which  will  liberate  these  sub- 
stances. The  reducing  gas  flame  forms  a  platinum  carbide,  so  care 
should  be  taken  to  heat  platinum  vessels  only  in  the  oxidizing  flame. 
Platinum  ware  may  be  cleansed  by  fusing  borax  in  it  and  it  may  be 
polished  by  rubbing  with  moist  sea-sand. 

Platinum  salts  are  unstable  and  readily  reduced,  liberating  the 
metal  as  a  black  porous  mass  called  "spongy  platinum"  or  a  black 
powder  called  "platinum  black." 

IMPORTANT  COMPOUNDS  OF  PLATINUM. 
Platinum  chloride,  PtCl4  +  5H2O. 
Chloroplatinic  acid,  H2PtCl6  +  6H2O. 


ARSENIC  39 

TESTS  FOR  PLATINUM. 

Use  a  separate  portion  of  a  solution  of  platinum  chloride,  PtCl4, 
for  each  of  the  following  tests: 

1.  Pass  H2S  gas  through  the  platinum  solution.    A  black  precipi- 
tate of  platinic  sulphide,  PtS2,  is  formed  slowly  in  a  cold  liquid, 
quickly  on  heating.    This  precipitate  is  insoluble  in  acids,  but  is 
soluble  in  yellow  ammonium  sulphide  T.S. 

2.  Add  NH4C1  T.S.  and  an  equal  volume  of  alcohol.    A  yellow 
precipitate  of  ammonium  chloroplatinate,  (NH4)2PtCl6  is  formed. 

Potassium  salts  give  a  yellow  precipitate  of  the  corresponding 
potassium  chloroplatinate,  K2PtCl6. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  PtCl4  +  H2S  = 

2.  PtCl4  +  NH4C1  = 

ARSENIC,  Asui  v  =  74.96. 

Arsenic  is  a  brittle  crystalline  element  which  has  the  physical 
properties  of  a  metal,  but  generally  behaves  as  a  non-metal.  The 
compounds  of  arsenic  are  dangerous  poisons  and  methods  for  the 
detection  of  small  quantities  of  the  element  in  the  stomach  contents 
and  in  various  tissues  of  the  body  are  of  great  importance  in  toxico- 
logical  chemistry. 

The  compounds  of  arsenic  are  generally  white,  though  several 
of  them  have  various  colors  due  to  the  other  elements  present. 
Most  of  the  important  compounds  of  arsenic  are  more  or  less  soluble 
in  water.  Arsenous  sulphide  and  Paris  green  are  insoluble.  Com- 
pounds of  arsenic  are  volatilized  at  high  temperatures,  with  or 
without  decomposition. 

IMPORTANT  COMPOUNDS  OF  ARSENIC. 

Arsenic  acid,  ortho-arsenic  acid,  (H3AsO4)2  +  H2O. 

Arsenous  iodide,  U.  S.  P.,  AsI3. 

Arsenous  oxide,  arsenic  trioxide,  U.  S.  P.,  arsenous  anhydride,  "  white 

arsenic,"  As2O3. 

Arsenous  sulphide,  "orpiment,"  As2S3. 

Copper  aceto-arsenite,  "  Paris  green,"  Cu(C2H3O2)2  +  3Cu(AsO2)2. 
Hydrogen  arsenide,  arseniuretted  hydrogen,  "arsine,"  AsH3. 
Potassium  arsenite,  K2HAsO3. 
Sodium  arsenate,  U.  S.  P.,  Na2HAsO4  +  7H2O. 

TESTS  FOR  ARSENIC. 

Use  a  separate  portion  of  a  solution  of  either  arsenous  chloride, 
AsCl3;  sodium  arsenite,  Na2HAsO3;  or  sodium  arsenate,  Na2HAsO4; 
for  tests  1,  2  and  3. 


40  THE  METALS 

1.  Pass  H2S  gas  through  the  warmed  arsenic  solution.    A  yellow 
precipitate  of  arsenous  sulphide,  As2S3,   is   formed,   quickly  from 
strong  solutions  of  arsenous  compounds  and   slowly  from  weak 
solutions  of  arsenous  compounds  or  from  solutions  of  arsenic  com- 
pounds. 

Collect  the  precipitate  on  a  filter  paper  and  pass  warm  (NH4)2S 
T.S.  through  the  paper.    The  precipitate  dissolves. 
Arsenous  sulphide  is  also  dissolved  by  (NH^COs  T.S. 

2.  REINSCH'S  TEST.— Mix  20  mils  of  HC1  T.S.  with  20  mils  of 
water  and  heat  to  boiling  in  a  small  porcelain  evaporating  dish, 
supported  on  a  piece  of  wire  gauze  on  a  tripod  over  the  flame  of  a 
Bunsen  burner.    When  the  liquid  boils  introduce  a  piece  of  bright 
copper  foil  about  2  cm.  square  and  continue  boiling  for  five  minutes, 
adding  water  from  time  to  time  to  keep  the  volume  of  the  liquid 
approximately   constant.    If  there  is  no  dark  deposit  formed  on 
the  copper,  the  reagent  and  apparatus  are  free  from  arsenic. 

A  dark  deposit  on  the  copper  shows  the  presence  of  arsenic  or 
other  easily  reduced  metal  in  the  reagents.  Such  a  preliminary  trial 
is  called  a  blank  test,  and  is  made  when  there  is  a  possibility  of  mis- 
leading results  being  obtained  from  the  reagents  or  other  source. 

If  the  copper  remains  bright  during  the  preliminary  test,  add  two 
or  three  drops  of  the  arsenic  solution  and  boil  for  ten  minutes,  adding 
water  from  time  to  time  to  keep  the  volume  constant.  A  dark 
coating  of  copper  arsenide,  Cu5As2,  will  be  deposited  on  the  copper, 
which  may  scale  off  if  much  arsenic  is  present.  Pour  off  the  liquid, 
rinse  the  copper,  first  with  water  and  then  with  alcohol,  and  then 
dry  it  by  pressing  between  the  folds  of  a  filter  paper.  Cut  the  copper 
into  small  pieces,  introduce  some  of  them  into  a  dry  test-tube,  and 
heat  the  lower  end  of  the  tube  by  holding  it  in  the  flame  of  a  Bunsen 
burner.  The  arsenic  will  oxidize  and  volatilize,  condensing  on  the 
cool  part  of  the  tube  as  a  white  deposit  of  arsenic  trioxide,  As2O3, 
whose  characteristic  octahedral  crystals  may  be  recognized  under  a 
magnifying  glass. 

Reinsch's  test  will  detect  very  small  amounts  of  arsenic,  and  it 
can  be  used  in  the  presence  of  organic  matter. 

3.  MARSH'S  TEST. — Prepare  a  small  hydrogen  generator  with  an 
outlet  tube  drawn  to  a  jet.    Introduce  a  few  pieces  of  pure  zinc  and 
some  diluted  hydrochloric  acid  into  the  generator  and  allow  the 
generated   hydrogen    to    escape   for   some   time,   to   displace   all 
of  the  air  in  the  generator.     To  test  this  hold  an  inverted  test- 
tube  down  over  the  jet  until  the  tube  is  filled  with  hydrogen. 
Remove  it  from  the  jet,  keeping  the  test-tube  inverted,  and  bring 
the  mouth  of  the  tube  in  contact  with  a  small  flame.    If  the  hydrogen 
burns  quietly  at  the  mouth  of  the  tube  it  is  safe  to  ignite  it  at  the 
jet,  but  not  otherwise.    To  make  it  doubly  sure  wrap  a  towel  loosely 


ANTIMONY  41 

around  the  flask,  when  the  hydrogen  burns  quietly  in  the  test-tube, 
and  then  bring  a  flame  to  the  jet  to  ignite  the  hydrogen.  The  towel 
can  then  be  removed.  Hold  a  cold  piece  of  porcelain  in  the  flame 
for  a  moment.  If  the  porcelain  is  not  discolored  the  reagents  and 
apparatus  are  free  from  arsenic.  If  a  dark  spot  appears  on  the 
porcelain,  the  reagents  or  apparatus  contain  arsenic  or  antimony. 

If  the  blank  test  shows  the  reagents  and  apparatus  to  be  free  of 
arsenic,  allow  the  hydrogen  to  continue  burning  and  introduce  a  few 
drops  of  an  arsenic  solution  into  the  generator  through  the  funnel 
tube.  The  arsenic  compound  is  reduced  by  a  small  part  of  the  nas- 
cent hydrogen  and  the  gaseous  compound  arsine,  AsH3,  is  formed 
which  mixes  with  the  rest  of  the  hydrogen.  In  a  short  time  the 
blue  hydrogen  flame  will  change  to  white,  the  arsine  burning  to 
form  arsenous  oxide,  As203,  and  water.  If  a  piece  of  cold  porcelain 
is  held  in  the  flame  it  lowers  the  temperature  of  the  burning  gas, 
and  a  lustrous  dark  spot  of  arsenic  will  be  deposited  on  the  porcelain. 
This  spot  will  dissolve  in  a  solution  of  chlorinated  soda,  NaClO. 

Marsh's  test  is  exceedingly  delicate,  but  if  a  mixture  containing 
organic  matter  is  to  be  tested  for  arsenic,  the  organic  matter  must  be 
destroyed  before  this  test  can  be  applied. 

4.  Add  magnesia  mixture  (see  reagents)  to  a  solution  of  Na2HAsO4. 
A  white  precipitate  of  magnesium  ammonium  arsenate,  MgNELiAsO^ 
will  form. 

Add  magnesia  mixture  to  a  solution  of  Na2HAsO3.  No  precipi- 
tate will  be  produced. 

This  test  will  help  to  distinguish  between  arsenous  and  arsenic 
compounds. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  AsCl3  +  H2S  = 
K2HAsO3  +  H2S  +  HC1  = 

2Na2HAs04  +  5H2S  +  4HC1  =  As2S3  +  4NaCl  +  S2  +  8H20 

2.  4AsCl3  +  lOCu  +  6H2O  =  2Cu6As2  +  12HC1  +  3O2 
2Cu5As2  +  3O2  =  10  Cu  +  2As2O3 

3.  AsCl3  +  Zn  +  HC1  = 
K2HAs03  +  Zn  +  HC1  = 

Na2HAsO4  +  4Zn  +  10HC1  =  2NaCl  +  4ZnCl2  +  4H2O  +AsH3 
As4  +  lONaCIO  +  6H2O  =  4H3As03  +  lONaCl 

4.  Na2HAs04  +  NH4OH  +  MgSO4  = 

ANTIMONY  (STIBIUM),  Sbui'v  =  120.2. 

Antimony  is  a  crystalline  brittle  metal  whose  compounds  are 
poisonous.  Type-metal  is  an  alloy  of  antimony  and  lead.  Soluble 
salts  of  antimony  are  decomposed  by  water  in  solutions  that  con- 


42  THE  METALS 

tain  moderate  amounts  of  mineral  acids  and  basic  compounds  con- 
taining the  monad  radical  antimony  I,  (SbO)',  or  more  complex 
oxides,  are  precipitated.  Organic  acids  or  large  amounts  of  mineral 
acids  prevent  this  precipitation. 

Antimony  salts  are  unstable  and  only  a  few  of  them  are  important. 
The  oxides  are  feebly  basic  with  acids  and  feebly  acid  with  bases. 
Antimonous  chloride  and  the  other  halogen  salts  are  deliquescent 
and  soluble  in  small  amounts  of  water  with  some  decomposition. 
Antimonyl-potassium  tartrate  is  moderately  soluble  in  water,  with- 
out decomposition. 

IMPORTANT  COMPOUNDS  OF  ANTIMONY. 

Antimonous  chloride,  "  butter  of  antimony,"  SbCl3. 

Antimonous  oxide,  Sb2O3. 

Antimonous  sulphide,  Sb2S3. 

Antimonyl-potassium  tartrate,  Antimony  and  potassium  tartrate, 

U.S.  P.,  "tartar  emetic,"  (KSbOC4H4O6)2H2O. 
Hydrogen  antimonide,  "stibine,"  SbH3. 

TESTS  FOR  ANTIMONY. 

Use  separate  portions  of  a  solution  of  antimonous  chloride,  SbCl3, 
for  each  of  the  following  tests : 

1 .  Pour  2  mils  of  the  antimony  chloride  solution  into  a  test-tube 
half  filled  with  water.     A  white  precipitate  consisting  mainly  of 
antimony  oxychloride,  SbOCl,  is  formed  if  the  solution  does  not 
contain  a  large  excess  of  free  acid. 

2.  Pass  H2S  gas  through  the  antimony  solution.    An  orange  red 
precipitate  of  antimony  sulphide,  Sb2S3,  is  formed.    Collect  a  small 
amount   of   the   precipitate   on   a   filter   and   pass   warm   yellow 
ammonium  sulphide  solution  through  the  filter.    The  precipitate 
dissolves. 

3.  Apply  Reinsch's  test,  including  the  blank  test,  as  described 
under  arsenic  on  page  4,  but  using  a  few  drops  of  the  antimony 
solution  instead  of  the  arsenic  solution.     A  dark  gray  deposit  of 
metallic  antimony  will  form  on  the  copper.    If  the  copper  with  the 
deposit  is  dried  and  cut  into  small  pieces  and  heated  in  a  dry  test- 
tube  the  antimony  will  oxidize  and  volatilize,  condensing  on  the 
cool  part  of  the  tube  as  a  white  deposit  of  antimonous  oxide,  Sb2O3, 
which  is  amorphous. 

4.  Apply  Marsh's  test,  including  the  blank  test,  as  described 
under  arsenic  on  pages  40-41,  but  using  a  few  drops  of  the  antimony 
solution,  instead  of  the  arsenic  solution.    The  antimony  compound 
is  reduced  by  a  small  part  of  the  nascent  hydrogen  and  gaseous 
stibine,  SbH3,  is  formed,  which  burns  with  the  excess  of  hydrogen 


TIN  43 

to  form  antimonous  oxide,  Sb2O3,  and  water.  If  cold  porcelain  is 
held  in  the  flame  the  antimony  will  be  deposited  on  it,  forming  dull, 
sooty  spots  which  will  not  dissolve  hi  a  solution  of  chlorinated  soda. 

Pass  the  gas  from  the  Marsh  generator  into  AgNO3  T.S.  If  stibine 
is  present  a  black  precipitate  of  silver  antimonide,  Ag3Sb,  is  formed. 
If  the  black  precipitate  appears  collect  it  on  a  filter  and  wash  well 
with  water.  Change  the  receiver  under  the  test-tube  and  pass 
diluted  HC1  through  the  filter,  refiltering  if  necessary  to  get  a  clear 
solution.  The  antimony  dissolves  as  SbCI3  and  the  silver  remains 
on  the  filter  as  AgCl.  Pass  H2S  through  the  filtrate  and  an  orange 
precipitate  of  Sb2S3  will  form. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  SbCI3  +  H20  = 

2.  SbCl3  +  H2S  = 

3.  Cu  +  SbCl3  = 
Sb  +  02  = 

4.  Zn  +  HC1  = 
SbCl3  +  H  = 
SbH3  +  02  = 
SbH8  +  AgNO,  = 
Ag3Sb  +  HC1  = 

TIN  (STANNtJM),  SniUv  =  119.0. 

Tin  is  a  soft,  white,  crystalline  metal  that  is  largely  used  in  the 
metallic  state  as  foil  and  pipe,  and  in  various  alloys  like  bronze, 
solder  and  pewter.  Tin  plate,  commonly  called  "tin,"  is  sheet  iron 
coated  with  tin. 

IMPORTANT  COMPOUNDS  OF  TIN. 

Stannous  chloride,  SnCl22H2O. 
Stannic  chloride,  SnCl45H2O. 
Stannous  sulphide,  SnS. 
Stannic  sulphide,  "Mosaic  gold,"  SnS2. 
Sodium  stannate,  Na2SnO34H2O. 

TESTS  FOR  TIN. 

Use  a  separate  portion  of  a  solution  of  stannous  chloride,  SnCl2, 
for  each  of  the  following  tests: 

1.  Pour  2  mils  of  the  stannous  chloride  solution  into  about  10 
mils  of  water.  A  precipitate  of  stannous  oxychloride,  Sn2OCl2,  will 
form,  if  the  solution  is  not  too  strongly  acid. 

Oxygen  from  the  air  frequently  precipitates  stannous  oxychloride 


44  THE  METALS 

with  formation  of  some  stannic  chloride  in  bottles  of  SnCl2  T.S. 
This  precipitation  is  retarded  by  acidulating  the  solution  with  HC1, 
and  keeping  some  pieces  of  metallic  tin  in  it. 

2.  Pass  H2S  through  the  SnCl2  solution.    A  brown  precipitate  of 
stannous  sulphide,  SnS,  is  formed.    Collect  a  small  amount  of  the 
precipitate  on  a  filter  and  pass  yellow  ammonium  sulphide  solution 
through  the  filter.    The  precipitate  will  dissolve,  forming  ammonium 
sulphostannate,  (NH4)2SnS3.    Add  an  excess  of  diluted  HC1  to  the 
filtrate.    A  yellow  precipitate  of  stannic  sulphide,  SnS2,  will  form. 

3.  Add  NaOH  T.S.  in  excess.    A  white  precipitate  of  stannous 
hydroxide,   Sn(OH)2,   appears  and  then  dissolves  in  the  excess  of 
the  reagent,  forming  sodium  stannite,  Na2SnO2.     Boil  the  solution 
gently.    A  white  precipitate  of  stannous  oxide,  SnO,  forms. 

4.  Add  a  few  drops  of  diluted  HOI  and  a  few  drops  of  HgCl2  T.S. 
to  the  solution  of  SnCl2.    A  white  precipitate  of  mercurous  chloride, 
HgCl,  will  form,  or,  if  the  SnCl2  is  in  large  excess,  the  precipitate 
will  be  gray  and  consist  of  metallic  mercury. 

5.  Pour  some  SnCl2  solution  into  a  test-tube  containing  zinc  and 
diluted  HC1.    Metallic  tin  is  liberated  as  a  gray,  spongy  precipitate 
which  will  form  SnCl2  and  dissolve  in  HC1  if  the  pieces  of  zinc  are 
removed . 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  SnCl2  +  H2O  = 

2.  SnCl2  +  H2S  = 
SnS  +  (NH4)2S2  = 
(NH4)2SnS3  +  HC1  =  SnS2  + 

3.  SnCl2  +  NaOH  = 
Sn(OH)2  +  NaOH  = 
NaaSnOa  +  H2O  = 

4.  SnCl2  +  HgCl2  =  HgCl  + 
SnCl2  +  HgCl2  =  Hg  + 

5.  SnCl2  +  Zn  = 
Sn  +  HC1  = 

NOTE  ON  STANNIC  COMPOUNDS. 

Stannic  compounds  in  solution  give  a  yellow  precipitate  of  stannic  sulphide, 
SnS2,  with  EbS,  and  a  white  precipitate  of  stannic  hydroxide,  Sn(OH)4,  with 
NaOH  T.S.,  soluble  in  an  excess  of  the  reagent.  Lead  and  hydrochloric  acid 
reduce  stannic  compounds,  SnCl2  being  formed,  which  may  be  recognized  by 
applying  tests  for  stannous  compounds  to  the  solution. 

SAMPLE  NUMBER  4. 
Analysis  of  a  solution  for  group  3  of  the  metals. 

The  solution  may  contain  salts  of  arsenic,  antimony,  tin,  gold 
or  platinum,  or  of  several  of  these  metals. 


TIN  45 

A.  Acidulate  10  mils  of  the  sample  with  HC1  T.S.,  warm  and 
pass  H2S  gas  through  it. 

If  no  precipitate  is  produced  on  passing  H2S  through  the  acid 
liquid  for  ten  minutes  or  more,  appreciable  amounts  of  the  metals 
of  group  3  are  absent. 

If  a  precipitate  forms,  continue  to  pass  H2S  through  the  liquid 
until  precipitation  appears  to  be  complete. 

Collect  the  precipitate  on  a  filter  and  test  the  filtrate  by  passing 
H2S  through  it  for  a  few  minutes.  If  no  more  precipitate  is  formed, 
group  3  is  all  precipitated,  and  the  filtrate  may  be  discarded,  unless 
it  is  to  be  kept  for  tests  for  succeeding  groups.  If  H2S  produces 
any  additional  precipitate,  continue  to  pass  the  gas  through  the 
filtrate  for  some  time  and  then  filter  through  the  same  filter  paper, 
testing  the  new  filtrate  in  the  same  manner  as  before. 

B.  Allow  the  water  to  drain  from  the  precipitate  on  the  filter 
paper.    Puncture  the  apex  of  the  filter  and  wash  the  precipitate  into 
a  test-tube  with  10  mils  of  (NH4)2CO8  T.S.,  which  will  dissolve 
As2S3,  but  not  the  sulphides  of  the  other  metals  of  this  group. 
Warm  the  liquid  by  standing  the  test-tube  in  hot  water  for  a  few 
minutes,  filter  and  wash  any  precipitate  on  the  filter  with  10  mils  of 
water,  discarding  the  washings. 

C.  Acidulate  the  filtrate  from  B  with  HC1  T.S.  and  pass  H2S 
through  it.    A  yellow  precipitate  shows  the  presence  of  arsenic.    A 
white  precipitate  consists  of  sulphur  and  may  be  disregarded.    If 
arsenic  is  found,  verify  it  by  applying  Reinsch's  test  to  the  liquid 
containing  the  precipitate. 

D.  Allow  any  precipitate  from  B  to  drain  on  the  filter  paper. 
Puncture  the  apex  of  the  paper  and  wash  the  residue  into  a  test-tube 
with  10  mils  of  concentrated  HC1.    Heat  the  test-tube  by  standing 
it  in  hot  water  for  ten  minutes  and  filter.    The  filtrate  will  contain 
the  antimony  and  tin,  if  present,  and  the  insoluble  residue  will  con- 
tain the  gold  and  platinum,  if  present. 

E.  Introduce  the  filtrate  from  D  into  a  small  generating  flask 
fitted  with  a  funnel  tube  and  a  delivery  tube.    Add  an  equal  volume 
of  water  and  a  small  amount  of  metallic  zinc.     Conduct  the  gas 
formed  through  a  half  test-tube  full  of  AgNO3  T.S.  for  three  minutes. 
If  no  precipitate  forms  in  the  AgNO3  T.S.,  antimony  is  absent.    A 
black  precipitate  indicates  the  presence  of  antimony. 

F.  If  a  black  precipitate  is  formed  in  the  AgNO3  T.S.  by  para- 
graph E,  collect  the  precipitate  on  a  filter  and  wash  with  water, 
rejecting  the  filtrate  and  washings.    Then  pass  HC1  T.S.  through 
the  filter,  passing  the  filtrate  through  the  filter  again  if  necessary 
to  get  a  clear  liquid.     Pass  H2S  through  the  filtrate.    An  orange- 
colored  precipitate  shows  the  presence  of  antimony. 

ON    Pass  a  small  amount  of  the  liquid  in  the  generator,  described 


46  THE  METALS 

* 

in  paragraph  E,  through  a  filter  into  about  1  mil  of  HgCl2  T.S. 
contained  in  a  test-tube.  A  white  or  gray  precipitate  shows  the 
presence  of  tin. 

//.  If  the  residue,  insoluble  in  hydrochloric  acid  is  light  colored 
the  sample  does  not  contain  gold  or  platinum  and  the  residue  may 
be  discarded.  If  this  residue  is  dark  colored  it  should  be  tested  for 
gold  and  platinum.  Puncture  the  apex  of  the  filter  paper  and  wash 
the  residue  into  a  test-tube  with  5  mils  of  nitrohydrochloric  acid. 
Warm  the  liquid  by  standing  the  test-tube  in  hot  water,  add  an 
equal  volume  of  water,  filter  out  any  insoluble  residue  and  divide 
the  clear  filtrate  into  two  parts. 

7.  To  one  part  of  the  filtrate  add  SnCl2  and  SnCl4  test  solutions 
and  let  stand.  A  purple  coloration  or  precipitate  indicates  gold. 
Verify  by  adding  SnCl2  and  SnCl4  test  solutions  to  the  original 
solution. 

J.  Neutralize  the  other  part  of  the  filtrate  from  H  with  XH4OH, 
add  an  equal  volume  of  alcohol  and  then  XH4C1  T.S.  A  yellow  or 
brown  precipitate  indicates  platinum.  Verify  by  adding  alcohol 
and  NH4C1  T.S.  to  the  original  solution. 

SAMPLE  NUMBER  o. 
Analysis  of  a  solution  for  groups  1,  2  and  3  of  the  metals. 

The  solution  should  be  neutral  or  acid  in  reaction.  If  alkaline,  add  HNOs 
to  the  portion  to  be  examined  until  the  reaction  is  acid. 

Add  HC1  as  long  as  a  precipitate  is  produced  and  filter. 
Precipitate— Pb,  Hgi,  Ag  Filtrate- 

Examine  for  metals  of    j      Warm  and  pass  H2S  as  long  as  a  precipitate  ie  pro- 
group  1  by  I.  duced.    Collect  the  ppt.  on  a  filter,  wash  and  drain. 
Digest  the  ppt.  with  (NH4)2Sx  and  filter. 

Residue—  Filtrate— 

Hg",  Pb,  Bi,  Cu,  Cd  As,  Sb,  Sn,  Au,  Pt 

Examine  for  metals  of        Examine  for  metals  of 
group  2  by  II.  group  3  by  III. 


I. 

Examination  of  any  precipitate  produced  by  HC1  in  a  neutral  or 
acid  solution. 

Wash  the  precipitate  on  the  filter  with  cold  water,  discarding  the  washings. 
Pass  a  portion  of  hot  water  through  the  washed  precipitate  on  the  filter. 

Filtrate— PbCl2  Residue— HgCl,  AgCl 

Add  K2CrO4.     Yellow      Pass  the  same  portion  of  XHiOH  through  the  residu  e 


ppt.,  PbCrCX,  shows 


lead 


on  the  filter  several  times 

Residue— XH2HgCl-j-Hg,    Filtrate— (XH»)8(AgCl), 
black,  shows  Add  an  excess  of  HXOs, 

white  ppt.,  AgCl,  shows 
mercurous  mercury  silver 


CALIFORNIA    COLLEGE 
Of    PHARMACY 

ANALYTICAL  TABLES 


47 


II. 

Examination  of  any  residue  insoluble  in  (XHt)^  from  a  precipi- 
tate produced  by  H2S  in  an  acid  liquid,  after  group  1  of  the  metals 
has  been  removed  from  a  solution. 


Wash  the  residue  on  the  filter,  first  with  an  additional  portion 
then  with  water,  and  drain,  discarding  the  washings. 

Pour  a  portion  of  hot  HXOs  through  the  filter  several  times. 
Residue-HgS  Filtrate—  Pb  (NO  ,)2>  Bi(NO,),,  Cu(NO,)2,  Cd(NOa)2 


x  and 


Black.  Dissolve  Add  cone.  HsSCh  and  boil  until  white  fumes  are  given  off,  cool, 
in  nitro- 
loric 
to 

Add  an  excess  of  XH^H  and  filter. 
White 


Prete~ 


expel  Cl,  dilute 
with  water,  fil- 
ter and  add 


dilute  with.  HzO  and  filter. 

Ffltrate-BMSOO.,  CuSO4,  CdSO, 


Dissolve  with 


SnCl2.    A  white    2 
or  grav  ppt. 
HgCf-Hg 
shows 


mercuric 
mercury 


and  add 


Yellow  ppt., 

PbCrO*, 

shows 


Precipitate, 
Bi(OHj, 

White.    Add 

XaOH  and 

SnCl2on  filter. 

Brown  or 
black  colora- 
tion,   BizOa, 
shows 


Filtrate— 


Divide  into  two  portions. 

Pass    HzS 
through  the 

other  portion, 
decolorizing 

with   KCX,   if 

Cu  is  present. 

A  yellow  ppt., 
CdS,  shows 


To  one  portion 

add  an  excess 

of   HdlLOj 

and  then 

K<Fe(CX)«. 

A  red  ppt.  or 

coloration, 

CuzFe(CN)., 

shows 


copper 


III. 


Examination  of  the  solution  hi  (XH^^x  of  any  precipitate  pro- 
duced by  HfjS  in  an  acid  liquid,  after  group  1  of  the  metals  has  been 
removed  from  a  solution. 


idulate  the  (XH4)jSx  solution  with  HC1  and  collect  any  precipitate  formed 
on  a  filter.  Transfer  the  precipitate  to  a  test  tube  and  digest  in  (XH4)jCOa  T.S. 
Filter. 

Filtrate—  Residue— Sb,S5,  SnS:,  Au.Sa,  PtS2 

NH,  ;AsS4 

Add  excess  of 


Wash  with  water.     Digest  in  concentrated  HC1.     Filter. 


HC1  and     ass 


Filtrate—  SbCls.  SnCli 


Residue—  Au2Sj,  PtS2 


H.S.  A  yellow    Place  in  H  generator.     Add    If    light    colored    Au    and    Pt 
ppt.,  AS&,       HiO  and  Zn.     Conduct  gat    are  absent.    If  dark  colored,  di- 
&hows  into  JfcXDnor  three  minutes  gest  in  nitrohydrochloric  acid. 


Add  HjO  and  filter. 


from  genera-    To  a  part  of 
tor  and  add  few  the  nitrate  add 
solve  in  HC1  drops  to  HgCl2.       SnCl2  and 
and  pass  HjS.      A  white  I  *  and  let 

An  orange         gray  ppt.,       stand.    A  pur- 
ple ppt.,  AujO, 
etc.,  shows 


ppt 
shows 


shows 


arsenic 


antimony 


Xeutralize 
another  part  of 
the  filtrate  with 
XH/JH,  add 

alcohol  and 
XH4C1.    Ayel- 


rm 


gold 


•howl 

platinum 


48  THE  METALS 

GROUP  4. 

Metals  precipitated   as  hydroxides   by   ammonium   hydroxide, 
NH4OH,  in  presence  of  ammonium  chloride,  NH4C1 : 
Iron,  Fe;  chromium,  Cr;  aluminum,  Al. 

IRON  (FERRUM)  Kl>™=  55.84 

Iron  is  largely  used  in  the  metallic  state,  nearly  pure  as  wrought 
iron,  less  pure  as  steel,  and  still  less  pure  as  cast  iron  and  malleable  iron. 
Reduced  iron  is  powdered  metallic  iron  obtained  by  the  action  of 
hydrogen  on  heated  ferric  oxide. 

In  the  ferrous  compounds,  iron  has  a  valence  of  two,  and  in  the 
ferric  compounds  it  has  a  valence  of  three.  Ferrous  compounds 
are  readily  oxidized  to  ferric  compounds  and  ferric  compounds  are 
nearly  as  readily  reduced  to  ferrous  compounds.  Both  series  of 
compounds  are  important.  Soluble  ferrous  compounds  are  generally 
light  green  in  color  and  slightly  color  their  solutions.  Soluble  ferric 
compounds  are  generally  brown  or  yellow  in  color  and  strongly 
color  their  solutions. 

IMPORTANT  FERROUS  COMPOUNDS. 

Ferrous  bromide,  FeBr26H2O. 

Ferrous  carbonate,  FeCO3. 

Ferrous  chloride,  FeCl2. 

Ferrous  hydroxide,  Fe(OH)2. 

Ferrous  iodide,  FeI2. 

Ferrous  oxide,  FeO. 

Ferrous  sulphate,  U.  S.  P.,  "  copperas,"  "  green  vitriol,"  FeSO47H2O. 

Exsiccated  ferrous  sulphate,  U.  S.  P.,  (FeSO4)23H2O. 

Ferrous  sulphide,  FeS. 

IMPORTANT  FERRIC  COMPOUNDS. 

Ferric  ammonium  sulphate,  "ferric  alum,"  NH4Fe(SO4)212H2O. 

Ferric  chloride,  U.  S.  P.,  FeCl36H2O. 

Ferric  ferrocyanide,  "Prussian  blue,"  Fe4(Fe(CN)6)3. 

Ferric  hydroxide,  "arsenic  antidote,"  Fe(OH)^. 

Ferric  hypophosphite,  N.  F.,  Fe(PH2O2)3. 

Ferric  nitrate,  Fe(NO3)3. 

Ferric  oxide,  "hematite,"  Fe2O3. 

Ferroso-ferric  oxide,  "magnetite,"  Fe3O4. 

Ferric  phosphate,  U.  S.  P.,  composition  indefinite. 

Ferric  pyrophosphate,  N.  F.,  composition  indefinite. 

Ferric  sulphate,  Fe2(SO4)3. 

Ferric  subsulphate,  "Monsel's  salt,"  Fe4O(SO4)5. 


IRON  49 

TESTS  FOR  FERROUS  IRON. 

Use  a  separate  portion  of  a  solution  of  a  ferrous  salt,  as  FeSCX,  for 
each  of  the  following  tests: 

1.  Add  NH4OH  T.S.  to  the  ferrous  solution.  A  white  precipi- 
tate of  ferrous  hydroxide,  Fe(OH)2,  is  produced  if  the  ferrous  salt  is 
pure,  but  the  precipitate  immediately  begins  to  oxidize  and  darken 
in  color,  forming  first  the  black  ferroso-ferric  hydroxide,  Fe3(OH)8, 
which  on  standing  changes  by  further  oxidation  to  the  light  brown 
ferric  hydroxide,  Fe(OH)3.  Ferrous  hydroxide  is  slightly  soluble  in 
solutions  of  ammonium  salts,  so  the  precipitation  of  iron  by  this 
reagent  is  not  complete.  NaOH  T.S.  will  produce  the  same  results. 

2. .  Add  (NH4)2CO3  T.S.  to  the  ferrous  solution.  A  white  precipi- 
tate of  ferrous  carbonate,  FeCO3,  is  produced  and  immediately  begins 
to  decompose  and  oxidize  and  darken  in  color,  forming  first  the 
black  ferroso-ferric  hydroxide,  Fe3(OH)8,  and  then,  by  further 
oxidation,  the  light  brown  ferric  hydroxide,  Fe(OH)3.  Na2CO3 
T.S.  will  produce  the  same  results. 

3.  Add  (XH4)2S  T.S.  to  the  ferrous  solution.    A  black  precipi- 
tate of  ferrous  sulphide,  FeS,  will  be  produced.     The  precipitate  is 
readily  soluble  in  acids. 

4.  'Add  K4Fe(CN)6  T.S.  to  the  ferrous  solution.    A  light  blue 
precipitate  of  potassium-ferrous  ferrocyanide,  K2FeFe(CN)6  will  be 
produced  which  will  readily  oxidize  to  form  ferric  ferrocyanide, 
Fe4(Fe(CN)6)3,  dark  blue.    The  precipitate  is  insoluble  in  acids. 
If  any  ferric  salt  is  present  in  the  ferrous  solution  the  precipitate 
will  be  dark  blue. 

5.  Add  K3Fe(CN)6  T.S.  to  the  ferrous  solution.    A  dark  blue 
precipitate  of  ferrous  ferricyanide,  Fe3(Fe(CN)6)2,  will  be  formed. 
The  precipitate  is  insoluble  in  acids. 

6.  Add  KCNS  T.S.  to  the  ferrous  solution.    No  coloration  or 
precipitation  will  be  produced,  unless  some  ferric  salt  is  present, 
which  would  give  a  bright  red  color  with  the  reagent. 

Complete  and  balance  the  following  equations: 

1.  FeS04  +  NH4OH  = 
FeCl2  +  NaOH  = 

2.  FeCl2  +  (NH4)2C03  - 
FeSO4  +  Na2CO3  = 

3.  FeSO4  +  (NH4)2S  = 
FeCl2  +  K2S  = 

4.  FeS04  +  K4Fe(CN)6  = 
FeBr2  +  K4Fe(CN)6  = 

5.  FeCl2  +  K3Fe(CN)6  = 
FeS04  +  K3Fe(CN)6  = 


50  THE  METALS 

TESTS  FOR  FERRIC  IRON. 

Use  a  separate  portion  of  a  solution  of  a  ferric  salt,  as  FeCl3,  for 
each  of  the  following  tests: 

1.  Add  NH4OH  T.S.  to  the  ferric  solution.    A  red-brown  gelati- 
nous precipitate  of  ferric  hydroxide,  Fe(OH)3,  will  be  formed,  which 
is  insoluble  in  an  excess  of  the  reagent  or  in  solutions  of  ammonium 
salts,  but  which  will  dissolve  in  acids.     Precipitation  is  not  complete 
if  glycerine,  sugar,  or  some  organic  acids  are  present.     On  boiling 
the  liquid  containing  the  precipitate,  it  changes  to  a  basic  ferric 
hydroxide,  FeO(OH),  which  is  not  gelatinous. 

NaOH  T.S.  produces  the  same  precipitate. 

2.  Add  Na^COs  T.S.  to  the  ferric  solution.     Carbon  dioxide  CO2, 
is  given  off,  and  ferric  hydroxide,  Fe(OH)3,  is  precipitated.    The 
same  result  is  obtained  with  (NH4)2CO3  T.S. 

3.  Add  (NH4)2S  T.S.  to  the  ferric  solution.    The  iron  is  reduced 
to  the  ferrous  condition  and  a  black  precipitate  is  obtained,  which 
consists  of  ferrous  sulphide,  FeS,  and  sulphur. 

H2S  reduces  ferric  salts  in  solution  to  the  ferrous  condition,  with 
precipitation  of  sulphur. 

4.  Add  K4Fe(CN)6  T.S.  to  the  ferric  solution.    A  deep  blue  pre- 
cipitate   of   ferric   ferrocyanide,    Fe4(Fe(CN)6)3,    is   formed.     The 
precipitate  is  insoluble  in  diluted  acids,  but  is  decomposed  by 
alkalis. 

5.  Add  K3Fe(CN)6  T.S.  to  the  ferric  solution.    Ferric  ferricya- 
nide,  FeFe(CN)6,  is  produced  and  changes  the  color  of  the  solution 
to  a  greenish-brown. 

6.  Add  KCNS  T.S.  to  the  ferric  solution.    A  blood  red  colora- 
tion is  produced,  due  to  the  formation  of  ferric  sulphocyanate, 
Fe(CNS)3.     This  reaction  is  most  dependable  in  the  presence  of 
moderate  amounts  of  free  HC1.   Many  different  compounds  prevent 
the  reaction  in  neutral  solutions. 

Complete  and  balance  the  following  equations: 

1.  FeCl3  +  NH4OH  = 
Fe2(S04)3  +  NaOH  = 
Fe(OH)3  +  heat  = 

2.  Fe(NO3)3  +  Na2CO3  +  H2O  = 
FeCl3  +  (NH4)2C03  +  H20  =  - 

3.  Fe2(SO4)3  +  3(NH4)2S  =  2FeS  +  S  +  3(NH4)2  S04 
Fe(N03)3  +  K2S  = 

FeCl3  +  H2S  = 

4.  FeCl3  +  K4Fe(CN)6  = 
Fe2(S04)3  +  K4Fe(CN)6  = 

5.  Fe(N03)3  +  K3Fe(CN)6  = 

6.  FeCl3  +  KCNS  = 
Fe3(S04)3  +  KCNS  = 


CHROMIUM  51 

CHROMIUM,  Cru'  Ui-  vi  =  52.0 

Chromium  is  a  light  gray  crystalline  metal  which  is  seldom  seen  in 
the  free  state.  It  is  contained  in  small  amounts  in  some  varieties  of 
steel,  to  which  it  imparts  additional  hardness. 

Compounds  of  chromium  are  brightly  colored.  When  it  has  a 
valence  of  two,  chromium  forms  the  unimportant  chromous  com- 
pounds; when  it  has  a  valence  of  three  it  forms  the  chromic  com- 
pounds; and  when  it  has  the  valence  of  six  it  is  an  acid-forming 
element,  in  chromic  acid  and  its  anhydride  and  the  chromates. 

IMPORTANT  COMPOUNDS  OF  CHROMIUM. 

Chromic  oxide,  Cr2O3. 

Chromium  trioxide,  U.  S.  P.,  chromic  anhydride,  "chromic  acid," 

Cr03. 

Chromium-potassium  sulphate,  "chrome  alum,"  KCr(SO4)212H2O. 
Chromium  sulphate,  Cr2(SO4)3H2O. 
Lead  chromate,  "chrome  yellow,"  PbCrO4. 
Potassium  chromate,  K2Cr04. 
Potassium  dichromate,  K2Cr2O7. 

TESTS  FOR  CHROMIC  SALTS. 

Use  a  separate  portion  of  a  solution  of  a  chromic  salt,  as  Cr2(SO4)3, 
for  each  of  tests,  1,  2,  3  and  4. 

1.  Add  NH4OH  T.S.  to  the  chromium  solution.    A  greenish-blue 
gelatinous  precipitate  of  chromic  hydroxide,  Cr(OH)3,  will  form. 

2.  Slowly   add   NaOH   T.S.   to   the   chromium   solution.    The 
greenish-blue  chromic  hydroxide,  Cr(OH)3,  will  form  and  re-dissolve 
in  an  excess  of  the  reagent  as  sodium  chromite,  NaCrO2.     Chromium 
hydroxide  is  re-precipitated  upon  boiling. 

3.  Add  Na2CO3  T.S.  to  the  chromium  solution.     The  greenish - 
blue  chromic  hydroxide,  Cr(OH)3,  will  be  precipitated  and  carbon 
dioxide  will  be  given  off. 

4.  Add  (NH4)2S  T.S.  to  the  chromium  solution.    The  greenish- 
blue  chromium  hydroxide,  Cr(OH)3,  will  be  precipitated  and  H2S 
will  be  given  off. 

5.  FUSION  TEST. — Collect  the  precipitate  from  any  of  the  preced- 
ing tests  for  chromium  on  a  filter  and  allow  to  drain.     Transfer  a 
portion  to  the  cover  of  a  porcelain  crucible,  add  about  an  equal 
amount  of  dry  KNO3  and  twice  as  much  dry  Na2CO3.     Support  the 
cover  on  a  pipe-stem  triangle  and  heat  in  the  Bunsen  burner  flame 
until  the  mass  is  fused.     The  KNO3  oxidizes  the  chromium  to  chro- 
mic acid,  which  forms  the  yellow  potassium  and  sodium  chromates, 

and  Na2CrO4.    Allow  the  mass  to  cool  and  dissolve  from  the 


52  THE  METALS 

cover  by  boiling  with  a  small  quantity  of  water  in  a  beaker.  Filte 
the  solution  formed,  acidulate  the  filtrate  with  HC2H3O2  and  ad 
Pb(C2H3O2)2  T.S.  A  yellow  precipitate  of  lead  chromate,  PbCrO 
will  be  formed. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  Cr2(S04)3  +  NH4OH  - 
CrCl3  +  NH4OH  = 

2.  Cr2(SO4)3  +  NaOH  = 
Cr(OH)3  +  NaOH  = 
NaCrO2  +  H2O  = 

3.  CrClg  +  Na2CO3  +  H2O  = 
Cr2(S04)3  +  (NH4)2C03  +  H2O  = 

4.  Cr2(S04)3  +  (NH4)2S  +  «2O  = 
CrCl3  +  (NH4)2S  +  H20  = 

5.  2O(OH)3  +  2KNO8  +  Na2CO3  =  K2Cr04  +  Na2Cr04  +  3H2( 

+  N202  +  CO2 

Cr(OH)3  +  KNO3  +  K2CO3  = 
K2Cr04  +  Pb(C2H302)2  = 

ALUMINUM,  Alui   =  27.1 

Aluminum  is  a  white  metal  which  is  largely  used  in  the  metalli 
state.  Its  compounds  are  generally  colorless. 

Aluminum  readily  forms  double  sulphates,  called  alums,  wit 
monad  metals  or  ammonium.  Similar  double  sulphates  of  oth( 
triad  and  monad  metals  are  also  called  alums. 

IMPORTANT  COMPOUNDS  OF  ALUMINUM. 

Aluminum-ammonium  sulphate,  alum,  U.  S.  P.,  ammonium  alun 

A1NH4(S04)2 12H20. 
Aluminum  chloride,  N.  F.,  A1C13. 
Aluminum  hydroxide,  U.  S.  P.,  A1(OH)3. 
Aluminum  oxide,  native  in  corundum,  emery,  etc.,  A1203. 
Aluminum-potassium  sulphate,  alum,  U.  S.  P.,  potassium  alun 

A1K(SO4)212H2O. 

Aluminum  silicates,  native  in  clays,  etc.,  composition  variable. 
Aluminum-sodium  sulphate,  sodium  alum,  AlNa(SO4)212H2O. 
Aluminum  sulphate,  N.  F.,  A12(SO4)316H2O. 
Exsiccated  alum,  U.  S.  P.,  "burnt  alum,"  "dried  alum,"  A1NH4(SO4 

or  A1K(S04)2. 

TESTS  FOR  ALUMINUM. 

Use  a  separate  portion  of  a  solution  of  an  aluminum  salt,  s 
A12(SO4)3,  for  each  of  the  following  tests: 


ALUMINUM  53 

1.  Add  NH4OH  T.S.  to  the  aluminum  solution  and  boil.    A 
white  gelatinous  precipitate  of  aluminum  hydroxide,  A1(OH)3  will 
form,  which  becomes  flocculent  upon  boiling. 

2.  Slowly  add  NaOH  T.S.  to  the  aluminum  solution.    A  white 
gelatinous  precipitate  of  aluminum  hydroxide,  A1(OH)3,  will  be 
produced  which  will  dissolve  in  an  excess  of  the  reagent  forming 
sodium  aluminate,  NaAlO2. 

3.  Add  Na2CO3  T.S.  to  the  aluminum  solution.    A  white  gelati- 
nous precipitate  of  aluminum  hydroxide,  A1(OH)3,  will  be  formed 
and  carbon  dioxide  will  be  liberated. 

4.  Add  (NH4)2S  T.S.  to  the  aluminum  solution.    The  white 
gelatinous  aluminum  hydroxide,  A1(OH)3,  will  be  precipitated  and 
H2S  will  be  generated. 

5.  Add  Na2HPO4  T.S.   to  the   aluminum   solution.    A  white 
precipitate  of  aluminum  phosphate,  A1PO4,  will  form.     This  precipi- 
tate is  soluble  in  mineral  acids  and  in  solutions  of  fixed  alkalies,  but 
is  insoluble  in  acetic  acid. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  A12(SO4)3  +  NH4OH  = 
A1C1,  +  NH4OH  = 

2.  A1C18  +  NaOH  = 
KA1(SO4)2  +  KOH  = 
A1(OH)8  +  NaOH  = 

3.  A12(S04)3  +  Na2CO3  +  H2O  = 
A1C13  +  (NH4)2C03  +  H20  = 

4.  A1C13  +  (NH4)2S  +  H20  = 
NH4A1(SO4)2  +  (NH4)2S  +  H2O  - 

5.  A12(S04)3  +  Na2HP04  = 
A1C18  +  Na2HPO4  = 

SAMPLE  NUMBER  6. 

Analysis  of  a  solution  for  group  4  of  the  metals. 

The  solution  may  contain  salts  of  aluminum,  chromium,  ferrous 
iron  or  ferric  iron. 

A.  To  about  10  mils  of  the  solution  add  a  few  drops  of  HNO3  T.S. 
and  boil  for  a  moment  to  oxidize  ferrous  iron,  if  present.  Then 
add  an  equal  volume  of  NH4C1  T.S.  and  an  excess  of  NH4OH  T.S.  to 
the  hot  liquid,  boil  again,  filter  and  wash  the  precipitate  on  the 
filter  paper.  The  filtrate  may  be  discarded  after  observing  that  it 
smells  of  ammonia,  which  shows  that  the  necessary  excess  of  this 
reagent  was  used. 


54  THE  METALS 

B.  Puncture  the  filter  and  transfer  the  precipitate  to  a  smal 
porcelain  evaporating  dish  with  about  5  mils  of  water.     Add  abou 
2  mils  of  NaOH  T.S.,  boil  for  several  minutes  and  filter,  if  necessary 

C.  Acidulate  the  filtrate  from  B  with  HC1  T.S.,  then  add  ai 
excess  of  (NH4)2CO3  T.S.  and  boil.    A  white  precipitate  shows  th< 
presence  of  aluminum. 

D.  Wash  the  precipitate  from  B  on  the  filter  paper  and  allow  i 
to  drain.     Transfer  a  part  of  the  moist  residue  to  the  cover  of  i 
porcelain  crucible,  cover  with  about  the  same  amount  of  dry  KNO 
and  about  twice  as  much  dry  Na2CO3,  place  the  cover  on  a  pipe 
stem  triangle  supported  on  a  tripod  and  heat  in  a  gas  flame  until  th< 
mass  is  fused.     Cool,  then  heat  the  cover  and  adhering  mass  witl 
about  10  mils  of  water  contained  in  a  small  beaker.    Remove  th< 
cover  from  the  liquid  and  filter. 

E.  Acidulate  the  filtrate  from  D  with  HC2H3O2  T.S.  and  adc 
Pb(C2H3O2)2  T.S.    A  yellow  precipitate   shows  the  presence   o 
chromium. 

F.  Wash  any  residue  from  D  on  the  filter  and  allow  it  to  drain 
Pass  about  5  mils  of  HC1  T.S.  through  the  filter  and  add  KCNS  T.S 
to  the  filtrate.     A  red  coloration  shows  the  presence  of  iron. 

G.  If  iron  is  found,  test  separate  portions  of  the  original  solutior 
for  ferrous  and  ferric  salts  as  follows : 

Acidulate  one  portion  with  HC1  T.S.  and  add  K8Fe(CN)6  T.S.  A 
dark  blue  precipitate  shows  the  presence  of  ferrous  iron. 

Acidulate  another  portion  with  HC1  T.S.  and  add  KCNS  T.S.  A 
red  coloration  shows  the  presence  of  ferric  iron. 

NOTE  ON  GROUP  4. 

Some  metals  of  the  fourth  and  fifth  groups  are  precipitated  with  the  fourtl 
group  if  either  phosphoric  acid  or  oxalic  acid  is  present,  as  their  phosphates  anc 
pxalates  are  insoluble  in  alkaline  liquids;  so,  in  the  separation  of  all  the  groups 
it  is  necessary  to  test  the  original  solution  for  phosphoric  acid  and  oxalic  acic 
and  remove  these  acids,  if  present,  by  methods  to  be  described  later,  before 
precipitating  Group  4. 


ANALYTICAL   TABLES 


55 


SAMPLE  NUMBER  7. 


Analysis  of  a  solution  for  groups  1,  2,  3  and  4  of  the  metals. 


The  solution  should  be  neutral  or  acid  in  reaction  to  litmus.  If  alkaline  add 
HNOs  to  the  portion  to  be  examined  until  the  reaction  is  acid.  Add  HC1  as  long 
as  a  precipitate  is  produced  and  filter. 


Precipitate 

Pb,  Hg', 

Ag. 

Examine  for 

metals  of  group  1 

by  I 


Filtrate 

Warm  and  pass  H2S  as  long  as  a  precipitate  is  produced. 

Filter. 
Precipitate 

Digest  with  (NH4)2SX  and  filter. 

Residue  Filtrate 

Hg",  Pb,  Sb,  Sn,  As, 

Bi,  Cu,  Cd.  Au,  Pt. 

Examine  for  Examine  for 

metals  of  group  2    metals  of  group  3 

by  II  by  III 


Filtrate, 

Boil  to  expel  H2S. 

Add  a  few  drops  of 

HNp3     and     boil 
again,  then  add 
NH4C1  and  an 

excess  of  NH4OH 

and  boil  again. 

Filter. 


Precipitate, 
Al,  Cr,  Fe. 

Examine  for 

metals  of  group  4 

by  IV 


I. 

Examination  of  any  precipitate  produced  by  HC1  in  a  neutral  or 
acid  solution. 


Wash  the  precipitate  on  the  filter  with  cold  water,  discarding  the  washings. 
Pass  a  portion  of  hot  water  through  the  washed  precipitate  on  the  filter. 


Filtrate— PbCl2 

Add  K2CrO4.     Yellow 
ppt.,  PbCrO4,  shows 


lead 


Residue— HgCl,  AgCl 

Pass  the  same  portion  of  NH4OH  through  the  residue 
on  the  filter  several  times 

Residue— NH2HgCl+Hg,     Filtrate,  (NH3)3(AgCl) 


black,  shows 


mercurous  mercury 


Add  an  excess  of  HNOa, 
white  ppt.,  AgCl,  shows 

silver 


56 


THE  METALS 


II. 

Examination  of  any  residue  insoluble  in  (NH4)2SX  from  a  precipi- 
tate produced  by  H2S  in  an  acid  liquid,  after  group  1  of  the  metals 
has  been  removed  from  a  solution. 

Wash  the  residue  on  the  filter,  first  with  an  additional  portion  of  (NH4)2SX  and 
then  with  water,  and  drain,  discarding  the  washings. 

Pour  a  portion  of  hot  HNO3  through  the  filter  several  times. 


Residue—  HgS 

Filtrate—  Pb(NO3)2,  Bi(NO3)3,  Cu(NO3)2,  Cd(NO3)2 

Black.  Dissolve 

Add  cone.  H2SO4  and  boil  until  white  fumes  are  given  off,  cool, 

in  nitro- 

dilute  with  H2O  and  filter. 

hydrochloric 

ppirl      ViOll    \~C\ 

Precipitate— 

Filtrate—  Bi2(SO4)3,  CuSO4,  CdSO4 

dL/lU.j    ULH1    IAJ 

expel  Cl,  dilute 

PbSO4 

Add  an  excess  of  NH4OH  and  filter. 

with  water,  fil- 
ter and  add 
SnCl2.    A  white 

White 
Dissolve  with 
NH4C2H302 

Precipitate, 
Bi(OH)3 

Filtrate— 
Cu(NH3)4SO4,  Cd(NH3)4SO4 

or  gray  ppt. 

and  add 

White.    Add 

Divide  into  two  portions. 

HgCl+Hg 
shows 

K2Cr2O7. 
Yellow  ppt., 
PbCrO4, 
shows 

NaOH  and 
SnCl2  on  filter. 
Brown  or 
black  colora- 
tion,  Bi2O3, 

To  one  portion 
add  an  excess 
of   HC2H3O2 
and  then 

K4Fe(CN)6. 

Pass    H2S 
through  the 
other  portion, 
decolorizing 
with   KCN,    if 

SI1O\VS 

A  red  ppt.  or 

Cu  is  present. 

coloration, 

A  yellow  ppt., 

Cu2Fe(CN)6, 

CdS,  shows 

mercuric 

shows 

mercury 

lead 

bismuth 

copper 

cadmium 

III. 

Examination  of  the  solution  in  (NH4)2SX  of  any  precipitate  pro- 
duced by  H2S  in  an  acid  liquid,  after  group  1  of  the  metals  has  been 
removed  from  a  solution. 

Acidulate  the  (NH4)2SX  solution  with  HC1  and  collect  any  precipitate  formed 
on  a  filter.  Transfer  the  precipitate  to  a  test  tube  and  digest  in  (NH4)2CO3  T.S. 
Filter. 


Filtrate  —      j                        Residue—  Sb2S5,  SnS2,  Au2S3,  PtS2 

(NH4)3AsS4 

Wash  with  water.     Digest  in  concentrated  HC1.     Filter. 

Add  excess  of 
HC1  and  pass 

Filtrate—  SbCls,  SnCl2 

Residue—  Au2S3,  PtS2 

H2S.  A  yellow 
ppt.,  As2S3, 

Place  in  H  generator.     Add 
H2O  and  Zn.     Conduct  gas 

If    light    colored    Au    and    Pt 
are  absent.    If  dark  colored,  di- 

shows 

into  AgNOs  for  three  minutes 

gest  in  nitrohydrochloric  acid. 

If  black  ppt. 

Filter    liquid 

Add  H2O  and  filter. 

forms  in 

from  genera- 

To  a  part  of 

Neutralize 

AgNO3  dis- 

tor and  add  few 

the  filtrate  add 

another  part  of 

solve  in  HC1 

drops  to  HgCl2. 

SnCl2  and 

the  filtrate  with 

and  pass  H2S. 

A  white  or 

SnCl4  and  let 

NH4OH,   add 

An  orange 
ppt.,  Sb2S3, 

gray  ppt., 
HgCl+Hg, 

stand.    A  pur- 
ple ppt.,  Au2O, 

alcohol  and 
NH4C1.    A  yel- 

shows 

shows 

etc.,  shows 

low  ppt., 

(NH4)2PtCl6, 

shows 

arsenic 

antimony 

tin 

gold 

platinum 

COBALT 


57 


IV. 

Examination  of  any  precipitate  produced  by  NH4OH,  after  groups 
1,  2  and  3  of  the  metals  have  been  removed  from  a  solution. 

Wash  the  precipitate  on  the  filter  with  water  and  drain,  discarding  the 
washings. 

Boil  the  precipitate  with  NaOH  for  several  minutes  and  filter. 


Filtrate,                                    Residue—  Cr(OH)3,  Fe(OH)3 

Na     U2.             Wash  with  H2O  and  drain,  discarding  the  washings.    Trans- 

Acidulate         fer  a  part  of  the  residue  to  a  porcelain  crucible  cover,  add  about 

with  HC1,  then 

the  same  amount  of  dry  KNO3  and  about  twice  as  much  dry 

add   an   excess 

Na2CO3  and  ignite  until  fused.    Cool,  heat  the  mass  with  H2O 

of  (NH4)2CO3  !  and  filter. 

and  boil.     A 
white  ppt., 

Filtrate, 

TT  PrH 

Residue  —  Fe2O3 

A1(OH)3, 

Aii\slN/4< 

Wash  on  the  filter  with  water,  discarding  the 

shows 

Acidulate 

washings.     Dissolve  in  HC1  and  add  KCNS. 

with 

A  red  coloration,  Fe(CNS)3,  shows 

HC2HSO2 

and  add 

iron 

Pb(C2H3O2)2. 

If  iron  is  present  test  separate  portions  of  the 

A  yellow 

original  solution  as  follows  : 

ppt.,  PbCrO4, 

To  one  portion  add  |  To  another  portion 

shows 

HC1  and  K3Fe(CN)6.     add  HC1  and  KCNS. 

A  dark  blue  ppt.,           A  red  coloration, 

Fe3(Fe(CN)6)2,  shows  |      Fe(CNS)3,  shows 

aluminum 

chromium 

ferrous  iron                    ferric  iron 

GROUP  5. 

Metals  precipitated  as  sulphides  by  ammonium  sulphide,  (NH^)2S, 
in  the  presence  of  ammonium  hydroxide,  NH4  OH: 
Cobalt,  Co;  nickel,  Ni;  manganese,  Mn;  zinc,  Zn. 


COBALT,  Co11'111  =  58.97 

Cobalt  is  a  light  colored  metal,  resembling  nickel,  which  is  seldom 
used  in  the  metallic  state.  Its  oxide  is  a  constituent  of  the  pigment 
"smalt,"  which  is  used  for  coloring  glass  blue.  The  cobaltic  com- 
pounds, in  which  the  metal  has  a  valence  of  three,  are  unimportant. 
In  the  more  important  cobaltous  salts  the  metal  has  a  valence  of  two. 
Soluble  cobalt  compounds  are  generally  red  in  color  when  hydrated 
or  in  solution  and  deep  blue  when  anhydrous. 

IMPORTANT  COMPOUNDS  OF  COBALT. 

Cobalt  chloride,  CoCl26H2O. 
Cobalt  nitrate,  Co(NO3)26H2O. 
Cobalt  oxide,  CoO. 
Cobalt  sulphate,  CoSO47H20. 


58  THE  METALS 

TESTS  FOR  COBALT. 

Use  a  separate  portion  of  a  solution  of  a  cobalt  salt,  as  cobalt 
nitrate,  Co(NO3)2,  for  each  of  the  following  tests: 

1.  Add  (NH4)2S  T.S.  to  the  cobalt  solution.     A  black  precipitate 
of  cobalt  sulphide,  CoS,  will  form,  which  is  insoluble  in  an  excess  of 
the  reagent  and  in  diluted  HC1,  but  is  soluble  in  concentrated  HC1 
and  in  diluted  HNO3. 

2.  Add  NaOH  T.S.  to  the  cobalt  solution  and  boil  the  liquid.     A 
blue  precipitate  of  a  basic  salt  is  obtained  which   forms  cobalt 
hydroxide,  Co(OH)2,  and  becomes  red  in  color  upon  boiling. 

3.  Slowly  add  NH4OH  T.S.  to  the  cobalt  solution.    A  precipi- 
tate of  cobalt  hydroxide,  Co(OH)2,  is  produced,  but  it  immediately 
dissolves  in  an  excess  of  the  reagent  to  form  a  brown  solution 
containing  complex  double  compounds  of  variable  composition. 

4.  Acidulate  a  portion  of  the  cobalt  solution  with  acetic  acid,  and 
add  twice  its  volume  of  KNO2  T.S.     Warm  the  mixture  and  allow  it 
to  stand  for  some  time.     A  yellow  crystalline  precipitate  of  potas- 
sium cobaltic  nitrite,  K3Co(NO2)e,  will  slowly  form.    This  com- 
pound is  soluble  in  water  but  is  insoluble  in  an  excess  of  KNO2. 

5.  Bend  the  end  of  a  piece  of  platinum  wire  around  the  point  of 
a  lead  pencil  to  form  a  loop  about  3  millimeters  in  diameter.    Heat 
this  loop  in  the  oxidizing  blowpipe  flame  for  a  moment  and  plunge 
it  into  some  powdered  borax.     Heat  the  adhering  borax  in  the 
blowpipe  flame  until  it  fuses  to  form  a  clear  and  colorless  glass-like 
bead.     Dip  this  bead  in  the  cobalt  solution  and  again  heat  it  in  the 
blowpipe  flame.     The  bead  will  be  colored  dark  blue.     If  too  much 
cobalt  is  present  the  color  will  appear  to  be  black. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  Co(N03)2  +  (NH4)2S  = 
CoS04  +  (NH4)2S  = 

2.  Co(NO3)2  +  NaOH  = 
CoCl2  +  KOH  = 

3.  CoCl2  +  NH4OH  = 
CoS04  +  NH4OH  = 

4.  CoCl2  +  7KNO2,  +  2HC2H302  =  K3Co(NO2)6  +  2KC1   + 

2KC2H302  +  H2O  +  NO 

NICKEL,  Nia'm=  58.68 

Nickel  is  a  silvery  white  metal  which  is  largely  used  in  the  metallic 
state  for  electroplating  and  as  a  constituent  of  various  alloys. 
German  silver  is  an  alloy  of  copper,  zinc  and  nickel.  United  States 
"nickel"  coins  are  75  per  cent,  copper  and  25  per  cent,  nickel. 


MANGANESE  59 

While  nickel  forms  two  series  of  compounds,  only  those  in  which  it 
has  a  valence  of  two  are  important.  Compounds  of  nickel  gener- 
ally have  a  bright  green  color  which  they  impart  to  their  solutions. 

IMPORTANT  COMPOUNDS  OF  NICKEL. 

Nickel-ammonium  sulphate,  Ni(NH4)2(SO4)2. 

Nickel  chloride,  NiCl26H2O. 

Nickel  oxide,  NiO. 

Nickel  sulphate,  NiSO47H2O. 

TESTS  FOR  NICKEL. 

Use  a  separate  portion  of  a  solution  of  a  nickel  salt,  as  nickel  chlo- 
ride, NiCl2,  for  each  of  the  following  tests: 

1.  Add  (NH4)2S  T.S.  to  the  nickel  solution.    A  black  precipitate 
of  nickel  sulphide,  NiS,  will  be  obtained.     This  precipitate  is  insol- 
uble in  acetic  acid  and  in  diluted  hydrochloric  acid,  but  is  readily 
soluble  in  nitric  acid  or  nitrohydrochloric  acid. 

2.  Add  NaOH  T.S.  to  the  nickel  solution.    A  pale  green  precipi- 
tate of  nickel  hydroxide,  Ni(OH)2,  will  be  formed,  which  is  insoluble 
in  an  excess  of  the  reagent. 

3.  Slowly  add  NH4OH  T.S.  to  the  nickel  solution.    The  pale 
green  precipitate  of  nickel  hydroxide,  Ni(OH)2,  will  be  formed  which 
will  form  a  double  salt  and  dissolve  to  a  blue  solution  in  an  excess 
of  the  reagent. 

4.  Make  a  borax  bead,  dip  it  in  the  nickel  solution  and  heat  it  in 
the  oxidizing  flame  of  the  blowpipe.     The  bead  will  be  colored  violet 
while  hot,  changing  to  brown  upon  cooling. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  NiCl2  +  (NH4)2S  = 
NiSO4  +  (NH4)2S  = 

2.  Ni(NO3)2  +  NaOH  = 
NiS04  +  KOH  = 

3.  NiCl,  +  NH4OH  = 
NiS04  +  NH4OH  = 

Ni(OH)2  +  2NH4C1  +  2NH4OH  =  (NH3)4NiCl2  +  4H2O. 

MANGANESE,  Mnu> Ui- iv-  vi'  vii  =  54.93 

Manganese  is  a  brittle  gray  metal,  resembling  cast  iron,  which  is 
seldom  seen  in  the  metallic  state.  It  is  a  constituent  of  some  varie- 
ties of  steel.  The  element  varies  widely  in  valence,  forming  several 
series  of  compounds,  of  which  the  most  important  are  the  manganous 


60  THE  METALS 

compounds  in  which  it  has  a  valence  of  two,  and  the  permanganates 
in  which  it  has  a  valence  of  seven. 

The  manganous  compounds  generally  have  a  light  pink  color. 
Other  compounds  of  manganese  have  various  bright  colors. 

IMPORTANT  COMPOUNDS  OF  MANGANESE. 

Manganese  chloride,  manganous  chloride,  MnCl24H2O. 
Manganese   dioxide,   precipitated   manganese   dioxide,    U.   S.    P., 

pyrolusite,  MnO2. 
Manganese     hypophosphite,     N.F.,    manganous    hypophosphite, 

Mn(PH202)2H20. 

Manganese  sulphate,  N.  F.,  manganous  sulphate,  MnS044H2O. 
Potassium  manganate,  K2Mn(>4. 
Potassium  permanganate,  KMn04. 

TESTS  FOR  MANGANOUS  SALTS. 

Use  a  separate  portion  of  a  solution  of  a  manganous  salt,  as 
MnCl2,  for  each  of  tests  1,  2,  3  and  4: 

1.  Add    (NH4)2S   T.S.   to   the  manganese   solution.    A   flesh- 
colored  precipitate  of  manganous  sulphide,  MnS,  will  be  formed. 
Manganous  sulphide  is  readily  soluble  in  diluted  mineral  acids  and 
in  acetic  acid. 

2.  Add  NaOH  T.S.  to  the  manganese  solution.    A  white  precipi- 
tate of  manganous  hydroxide,  Mn(OH)2,  will  be  formed,  which 
oxidizes  to  form  basic  manganic  hydroxide,  MnO(OH),  and  turns 
brown  upon  exposure  to  air.    Both  hydroxides  are  insoluble  in  an 
excess  of  the  reagent. 

Collect  some  of  the  precipitated  Mn(OH)2  on  a  filter  paper  for  use 
in  tests  5  and  6. 

3.  Add  NH4OH  T.S.  to  the  manganese  solution.    A  part  of  the 
manganese   is  precipitated   as   the   white  manganous  hydroxide, 
Mn(OH)2,  which  oxidizes  and  turns  brown  upon  exposure  to  air. 
The  other  part  forms  a  double  salt  and  remains  in  solution. 

To  a  portion  of  the  manganese  solution  add  double  its  volume  of 
NH4C1  T.S.  and  then  add  NH4OH  T.S.  The  manganese  forms  a 
soluble  double  salt,  manganous  ammonium  chloride,  Mn(XH4)2Cl4, 
and  precipitation  does  not  occur  until  oxidation  takes  place  when  the 
brown  basic  manganic  hydroxide,  MnO(OH),  is  slowly  precipitated. 

4.  Add  Na2CO3  T.S.  to  the  manganese  solution.    A  white  pre- 
cipitate of  manganous  carbonate,  MnCO3,  will  form,  which  oxidizes 
to  form  the  brown  basic  manganic  hydroxide,  MnO(OH),  upon 
exposure  to  air. 

5.  Mix  about  one-quarter  of  a  teaspoonful  of  lead  peroxide,  Pb02, 
with  about  5  mils  of  HNO3  T.S.  and  a  small  quantity  of  the  pre- 


ZINC  61 

cipitated  Mn(OH)2  obtained  in  test  2,  and  boil.  Allow  the  powder 
to  settle  and  observe  that  the  supernatant  liquid  has  the  pink  or 
purplish  color  of  permanganic  acid,  HMnO4.  The  Mn(OH)2  dis- 
solves to  form  Mn(NO3)2,  which  is  then  oxidized  to  HMnO4. 

6»  FUSION  TEST.— Place  some  of  the  precipitated  Mn(OH)2 
from  test  2  on  the  cover  of  a  porcelain  crucible  and  cover  it  with 
about  the  same  amount  of  dry  KNO3  and  twice  as  much  dry  Na2CO3. 
Support  the  cover  on  a  pipe-stem  triangle  and  heat  in  the  flame  of 
a  Bunsen  burner  until  fused.  Allow  to  cool.  The  manganese  is 
oxidized  to  manganic  acid  by  the  potassium  nitrate  and  a  green 
mass  containing  potassium  and  sodium  manganates  will  be  obtained. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  MnCl2  +  (NH4)2S  = 
MnSO4  4-  (NH4)2S  = 

2.  MnSO4  +  NaOH  = 
MnCl2  +  KOH  = 

4Mn(OH)2  +  O2  =  4MnO(OH)  +  2H20 

3.  2MnSO4  +  2NH4OH  =  Mn(NH4)2(S04)2  +  Mn(OH)2 
MnCl2  +  2NH4C1  =  Mn(NH4)2Cl4 

4.  Mn(XO3)2  +  Na2CO3  = 
MnS04  +  (NH4)2C03  = 

5.  2Mn(NO3)2  +  5PbO2  +  6HNO3  =  2HMnO4  +  5Pb(NO3)2  + 

2H2O 

6.  3Mn(OH)2  +  4KNO3  +  Na2CO3  =  2K2MnO4  +  Na2MnO4  + 

4NO  +  CO2  +  3H2O 

ZINC,  Zn11    =  65.37 

Zinc  is  a  bluish-white  crystalline  metal  which  is  largely  used  in  the 
metallic  state.  Brass  is  an  alloy  of  copper  and  zinc.  Galvanized 
iron  is  sheet  iron  coated  with  zinc.  Zinc  salts  are  poisonous  and  are 
generally  colorless. 

IMPORTANT  COMPOUNDS  OF  ZINC. 

Zinc  acetate,  U.  S.  P.,  Zn(C2H3O2)2. 

Zinc  bromide,  ZnBr2. 

Zinc  sub-carbonate,  precipitated  zinc  carbonate,  U.  S.  P.,  basic  zinc 

carbonate,  (ZnCO3)2(Zn(OH)2)3. 
Zinc  chloride,  U.  S.  P.,  ZnCl2. 
Zinc  iodide,  ZnI2. 

Zinc  oxide,  U.  S.  P.,  "zinc  white,"  ZnO. 
Zinc  sulphate,  U.  S.  P.,  "white  vitriol/'  ZnSO47H2O. 


62  .  THE  METALS 

TESTS  FOE  ZINC. 

Use  a  separate  portion  of  a  solution  of  a  zinc  salt,  as  ZnSO4,  for 
each  of  tests  1,  2,  3  and  4. 

1.  Add  (NH4)2S  T.S.  to  the  zinc  solution.     A  white  p  ecipitate  of 
zinc  sulphide,  ZnS,  will  be  obtained.     The  precipitate  is  soluble  in 
mineral  acids,  but  is  not  soluble  in  acetic  acid. 

2.  Slowly  add  NaOH  T.S.  to  the  zinc  solution.    A  white  precipi- 
tate of  zinc  hydroxide,  Zn(OH)2  is  obtained,  which  forms  sodium 
zincate,  Na2ZnO2,  and  dissolves  in  an  excess  of  the  reagent.    The 
zinc  is  partially  re-precipitated  as  zinc  oxide,  ZnO,  upon  boiling  the 
solution. 

3.  Slowly  add  NH4OH  T.S.  to  the  zinc  solution  until  the  precipi- 
tate formed  is  re-dissolved,  then  boil  the  solution  obtained.    The  first 
precipitate  formed  is  zinc  hydroxide,  Zn(OH)2.     This  dissolves  in  an 
excess  of  NH4OH  to  form  ammonium  zincate,  (NH4)2ZnO2.     Finally 
a  precipitate  of  zinc  oxide,  ZnO,  is  formed  upon  boiling  the  solution. 

4.  Add  Na2CO3  T.S.  to  the  zinc  solution.    A  white  precipitate  of 
zinc  sub-carbonate,  (ZnCO3)2  (Zn(OH)2)3,  is  obtained.    This  precipi- 
tate is  soluble  in  acids  and  in  alkali  hydroxides. 

Collect  some  of  the  precipitated  zinc  sub-carbonate  on  a  small 
filter  paper,  wash  with  water  and  allow  to  drain  for  use  in  test  5. 

5.  Fold  the  filter  paper  containing  the  precipitate  collected  from 
test  4,  place  it  on  the  cover  of  a  porcelain  crucible  with  a  small 
amount  of  dry  Na2CO3  and  add  a  few  drops  of  Co(NO3)2  T.S. 
Place  the  cover  on  a  pipe-stem  triangle  supported  on  a  tripod  and 
ignite  in  the  flame  of  a  gas  burner.    A  green  mass  consisting  of  zinc 
and  cobalt  oxides  will  be  obtained. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  ZnSO4  +  (NH4)2S  = 
Zn(C2H302)2  +  K2S  = 

2.  ZnCl2  +  NaOH  = 
Zn(C2H302)2  +  KOH  = 
Zn(OH)2  +  NaOH  = 

3.  ZnSO4  +  NH4OH  - 
Zn(OH)2  +  NH4OH  = 
(NH4)2ZnO2  +  heat  = 

4.  5ZnSO4  +  5Na2CO3  +  3H2O  =  (ZnCO3)2(Zn(OH)2)3  +  5Na2SO4 

+  3C02 
ZnCl2  +  K2CO3  = 


ZINC  63 

SAMPLE  NUMBER  8. 
Analysis  of  a  solution  for  group  5  of  the  metals. 

The  sample  may  contain  salts  of  any  of  the  following  metals: 
cobalt,  nickel,  manganese,  zinc. 

A.  To  about  10  mils  of  the  solution  add  (NH4)2S  T.S.     If  no 
precipitate  is  produced,  the  metals  of  group  5  are  absent.     If  a  pre- 
cipitate is  produced,  continue  to  add  the  reagent  until  precipitation 
appears  to  be  complete.     Filter  and  test  the  filtrate.     If  precipita- 
tion is  complete,  the  filtrate  may  be  discarded. 

B.  Wash  the  precipitate  on  the  filter  with  cold  water,  discarding 
the  washings,  then  pass  the  same  portion  of  about  5  mils  of  HC1 
T.S.  through  the  precipitate  on  the  filter  several  times.     Cobalt 
and  nickel  sulphides,  if  present,  remain  undissolved,  and  zinc  and 
manganese,  if  present,  form  chlorides  and  dissolve. 

C.  Boil  the  filtrate  from  B  until  it  ceases  to  smell  of  H2S.    Add 
an  excess  of  NaOH  T.S.  and  filter,  if  necessary.     Manganese,  if 
present,  is  precipitated  as  Mn(OH)2,  and  zinc,  if  present,  forms 
Na2ZnO2  and  remains  in  solution. 

D.  Pass  H2S  through  the  filtrate  from  C.     If  no  precipitate  is 
produced,  zinc  is  absent.     A  white  or  gray  precipitate,  ZnS,  shows 
the  presence  of  zinc. 

E.  If  NaOH  produced  no  precipitate  in  C,  manganese  is  absent. 
If  a  precipitate  was  produced  by  NaOH  in  C,  transfer  some  of  it 
from  the  filter  to  a  test-tube,  add  a  small   amount  of  powdered 
lead  peroxide,  PbO2,  and  about  5  mils  of  HNO3  T.S.,  and  boil. 
Allow  the  powder  to  settle.     If  the  supernatant  liquid  has  a  pink 
color,  NaMnO4,  the  presence  of  manganese  is  shown. 

F.  If  there  is  no  dark-colored  residue  on  the  filter  paper  from  B, 
cobalt  and  nickel  are  absent.     If  there  is  a  dark-colored  residue 
insoluble  in  HC1  T.S.,  in  B,  puncture  the  filter  and  wash  the  residue 
into  a  test-tube  with  about  5  mils  of  nitro-hydrochloric  acid.     Add 
an  excess  of  NaOH  T.S.  and  collect  the  precipitate  produced  on  a 
filter  and  wash  with  water,  discarding  the  filtrate  and  washings. 
Pass  the  same  portion  of  about  5  mils  of  HC2H3O2  T.S.  through  the 
precipitate  on  .the  filter  several  times. 

G.  To  the  solution  in  HC2H302,  from  F,  add  about  10  mils  each 
of  KC1  T.S.  and  KNO2  T.S.,  let  stand  for  half  an  hour  and  filter  if 
necessary.     If  no  precipitate  is  produced  cobalt  is  absent.     If  KC1 
and  KNO2  produce  a  yellow  precipitate,  K3Co(NO2)6,  the  presence 
of  cobalt  is  shown. 

H.  To  the  filtrate  from  G,  add  an  excess  of  NaOH.  If  no 
precipitate  is  produced,  nickel  is  absent.  A  light  green  precipitate 
Ni(OH)2,*shows  the  presence  of  nickel. 


64 


THE  METALS 


SAMPLE  NUMBER  9. 
Analysis  of  a  solution  for  groups  1,  2,  3,  4  and  5  of  the  metals, 


The  solution  should  be  neutral  or  acid  in  reaction  to  litmus.  If  alkaline  adc 
HNOs  to  the  portion  to  be  examined  for  metals  until  the  reaction  is  acid.  Adc 
HC1  as  long  as  a  precipitate  is  produced  and  filter. 


Precipitate 

Filtrate 

Pb,  Hg', 
Ag. 

Warm  and  pass  H2S  as  long  as  a  precipitate  is  produced.    Filter 

Examine  for 

Precipitate 

Filtrate 

metals  of 

Digest  with  (NH4)2SX  and 

Boil  to  expel  H2S.    Add  a  fev 

group  1  by 

filter. 

Residue             Filtrate 
Hg",  Pb,          Sb,  Sn,  As, 

drops  of  HNO3  and  boil  again 
then  add  NH4C1  and  an  exces 
of  NH4OH  and  boil  again.  Filter 

Bi,  Cu,  Cd.    |       Au,  Pt. 

Precipitate 

Filtrate 

Examine 

Examine 

Al,  Cr,  Fe. 

Pass   H2S    as 

for  metals  of 
group  2  by 

for  metals  of 
group  3  by 

Examine 
for  metals  of 
group  4  by 
IV  . 

long  as  a  ppt 
is  produced. 
Filter. 

Precipitate 

Zn,  Mn, 

Co,  Ni. 

Examine 

for   metals   of 

group  5  by 

I. 

Examination  of  any  precipitate  produced  by  HC1  in  a  neutral  o 
acid  solution. 


Wash  the  precipitate  on  the  filter  with  cold  water,  discarding  the  washings. 
Pass  a  portion  of  hot  water  through  the  washed  precipitate  on  the  filter. 


Filtrate— PbCl2 

Add  H2SO4  and  cool. 
White  ppt.,  PbSO4,  shows 


Residue— HgCl,  AgCl 

Pass  the  same  portion  of  NH4OH  through  the  residui 
on  the  filter  several  times 


lead 

Residue—  NH2HgCl+Hg, 

black,  shows 

mercurous  mercury 

Filtrate,  (NH3)3(AgCl)2. 

Add  an  excess  of  HNO3 
white  ppt.,  AgCl,  show 

silver 

ANALYTICAL  TABLES 


65 


II. 

Examination  of  any  residue  insoluble  in  (NH4)2SX  from  a  precipi- 
tate produced  by  H2S  in  an  acid  liquid,  after  group  1  of  the  metals 
has  been  removed  from  a  solution. 

Wash  the  residue  on  the  filter,  first  with  an  additional  portion  of  (NH4)2SX  and 
then  with  water,  and  drain,  discarding  the  washings. 

Pour  a  portion  of  hot  HNOs  through  the  filter  several  times. 


Residue—  HgS 

Filtrate—  Pb(N03)2,  Bi(NO3)3,  Cu(NO3)2,  Cd(NO3)2 

Black.  Dissolve 

Add  cone.  H2SO4  and  boil  until  white  fumes  are  given  off,  cool, 

in  nitro- 

dilute  with  H2O  and  filter. 

hydro  chloric 
acid,  boil  to 

Precipitate  — 

^\l~C9f\ 

Filtrate—  Bi2(SO4)3,  CuSO4,  CdSO4 

expel  Cl,  dilute 

FbSO4 

Add  an  excess  of  NH4OH  and  filter. 

with  water,  fil- 
ter and  add 
SnCl2.    A  white 

White 
Dissolve  with 
NH4C2H3O2 

Precipitate, 
Bi(OH)3 

Filtrate— 
Cu(NH3)4SO4,  Cd(NH3)4SO4 

°Hggay+&pgt- 

shows 

and  add 
K2Cr2O7. 
Yellow  ppt., 
PbCr04, 
shows 

White.    Add 
NaOH  and 
SnCl2  on  filter. 
Brown  or 
black  colora- 
tion,  Bi2O3, 

Divide  into 

To  one  portion 
add  an  excess 
of   HC2H3O2 
and  then 

K4Fe(CN)6. 

two  portions. 

Pass    H2S 
through  the 
other  portion, 
decolorizing 
with   KCN,    if 

SI1O\VS 

A  red  ppt.  or 

Cu  is  present. 

coloration, 

Cu2Fe(CN)6, 

A  yellow  ppt.,  • 
CdS,  shows 

mercuric 

shows 

mercury 

lead 

bismuth 

copper 

cadmium 

III. 

Examination  of  the  solution  in  (NH4)2SX  of  any  precipitate  pro- 
duced by  H2S  in  an  acid  liquid,  after  group  1  of  the  metals  has  been 
removed  from  a  solution. 

Acidulate  the  (NH4)2SX  solution  with  HC1  and  collect  any  precipitate  formed 
on  a  filter.  Transfer  the  precipitate  to  a  test  tube  and  digest  in  (NH4)2CO3  T.S. 
Filter. 


Filtrate—                             Residue—  Sb2S5,  SnS2,  Au2S3,  PtS2 

(NH4)3AsS4 
Wash  with  water.     Digest  in  concentrated 

HC1.     Filter. 

Arlrl     AYr»PQQ    r\f   • 

<rxLHJ      t^AL-t/Oo     Ul    ,                TP-il-f-ro-fA         C 

HC1  and  pass 

bC!3,  SnCl2 

Residue—  Au2S3,  PtS2 

H2S.  A  yellow    Place  in  H  generator.     Add 

If    light    colored    Au    and    Pt 

ppt.,  As2S3,       H2O  and  Zn. 

Conduct  gas 

are  absent.    If 

dark  colored,  di- 

shows           into  AgNOs  for  three  minutes 

gest  in  nitrohydrochloric  acid. 

If  black  ppt. 

Filter    liquid 

Add  H2O  and  filter. 

forms  in 

from  genera- 

To a  part  of 

Neutralize 

AgNOs  dis- 

tor and  add  few 

the  filtrate  add 

another  part  of 

solve  in  HC1 

drops  to  HgCl2. 

SnCl2  and 

the  filtrate  with 

and  pass  H2S. 

A  white  or 

SnCl4  and  let 

NHUOH,  add 

An  orange 

gray  ppt., 

stand.    A  pur- 

alcohol and 

ppt.,  Sb2S3, 

HgCl4-Hg, 

ple  ppt.,  Au2O, 

NH4C1.    A  yel- 

shows 

shows 

etc.,  shows 

low  ppt., 
(NH4)2PtCl6, 

shows 

arsenic 

antimony 

tin 

gold 

platinum 

66 


THE  METALS 


IV. 

Examination  of  any  precipitate  produced  by  NH4OH,  after  groups 
1,  2  and  3  of  the  metals  have  been  removed  from  a  solution. 


Wash  the  precipitate  on  the  filter  with  water  and  drain,  discarding  the 
washings. 

Boil  the  precipitate  with  NaOH  for  several  minutes  and  filter. 


Filtrate, 

Residue  —  Cr(OH)3,  Fe(OH)3 

NaAlO2. 

Wash  with  H2O  and  drain,  discarding  the  washings.    Trans- 

Acidulate 

fer  a  part  of  the  residue  to  a  porcelain  crucible  cover,  add  aboul 

with  HC1, 

the  same  amount  of  dry  KNO3  and  about  twice  as  much  drj 

then  add  an 

Na2CO3  and  ignite  until  fused.    Cool,  heat  the  mass  with  H2C 

excess  of 

and  filter. 

(NH4)2C03 
and  boil. 

Filtrate, 

Krt~r\ 

Residue 

—  Fe2O3 

A  white  ppt., 

2orU4. 

Wash  on  the  filter  with  water,  discarding  the 

A1(OH)3, 

Acidulate 

washings.    Dissolve  in  HC1  and  add  KCNS 

shows 

with 

A  red  coloration,  Fe(CNS)3,  shows 

HC2H3O2 
and  add 

iron 

Pb(C2H302)2. 

If  iron  is  present  test  separate  portions  of  the 

A  yellow 

original  solution  as  follows: 

ppt.,  PbCr04, 
shows 

To  one  portion  add 
HC1  and  K3Fe(CN)6. 

To  another  portion 
add  HC1  and  KCNS. 

A  dark  blue  ppt., 

A  red  coloration, 

Fe3(Fe(CN)6)2l  shows 

Fe(CNS)3,  shows 

aluminum 

chromium 

ferrous  iron 

ferric  iron 

V. 

Examination  of  any  precipitate  produced  by  H2S  in  an  alkaline 
liquid,  after  groups  1,  2,  3  and  4  of  the  metals  have  been  removed 
from  a  solution. 


Wash  the  precipitate  on  the  filter  with  water  and  drain,  discarding  the  wash- 
ings.    Pass  the  same  portion  of  HC1  through  the  filter  several  times. 


Filtrate—  ZnCl2,  MnCl2. 

Residue  —  CoS,  NiS. 

Boil  to  expel  H2S,  cool,  add 

If  light  colored,  cobalt  and  nickel  are  absent 

excess  of  NaOH  and  filter. 

If  dark  colored,  dissolve  in  nitro-hydrochloric 

Filtrate, 
Na2ZnO2. 

Residue, 
Mn(OH)2. 

acid,  add  an  excess  of  NaOH,  collect  the  ppt 
produced  on  a  filter  and  wash,  discarding  the 
washings.    Dissolve  the  ppt.  in  HC2H302,  adc 

Pass  H2S. 

Boil  with  cone. 

an  excess  of  KC1  and  of  KNO2,  let  stand  foi 

A  white  ppt. 

HNO3  and 

half  an  hour  and  filter. 

ZnS,  shows 

PbO2   and  let 

stand.   A  pink 

Precipitate, 

Filtrate,  Ni(C2H3O2)2 

supernatant 
liquid  shows 

K3Co(N02)6, 
yellow,  shows 

Add  an  excess  of 
NaOH.    Light  green 

ppt.  Ni(OH)2,  shows 

zinc 

manganese 

cobalt 

nickel 

CALCIUM  67 

GROUP  6. 

Metals  precipitated  as  carbonates  from  a  neutral  or  alkaline 
solution  by  ammonium  carbonate,  (NH4)2CO3  in  the  presence  of 
ammonium  chloride,  NH4C1: 

Calcium,  Ca;  strontium,  Sr;  barium,  Ba. 

CALCIUM,  Ca"  =  40.07 

Calcium  is  a  light  colored  metal  which  quickly  oxidizes  in  the 
air;  it  is  seldom  seen  in  the  metallic  state,  but  many  of  its  com- 
pounds are  important.  Many  of  the  compounds  of  calcium  were 
formerly  named  as  salts  of  lime,  for  example;  calcium  phosphate  was 
formerly  called  "phosphate  of  lime." 

IMPORTANT  COMPOUNDS  OF  CALCIUM. 

Calcium  acetate,  Ca(C2H3O2)2. 

Calcium  bromide,  U.  S.  P.,  CaBr2. 

Calcium  carbonate,  precipitated  calcium  carbonate,  U.  S.  P.,  pre- 
pared chalk,  U.  S.  P.,  marble,  chalk,  limestone,  whiting,  CaCO3. 

Calcium  carbide,  CaC2. 

Calcium  chloride,  U.  S.  P.,  CaCl2. 

Calcium  hydroxide,  slaked  lime,  Ca(OH)2. 

Calcium  hypophosphite,  U.  S.  P.,  Ca(PH2O2)2. 

Calcium  iodide,  CaI2. 

Calcium  oxide,  U.  S.  P.,  lime,  quicklime,  CaO. 

Calcium  phosphate,  precipitated  calcium  phosphate,  N.F.,  Ca3(PO4)2. 

Calcium  sulphate,  gypsum,  CaSO42H2O. 

Exsiccated  calcium  sulphate,  plaster  of  Paris,  CaSO4. 

Calcium  sulphide,  CaS. 

Calcium  sulphite,  CaSO32H2O. 

Chlorinated  lime,  U.  S.  P.,  bleaching  powder,  "chloride  of  lime/' 
CaOCl2. 

TESTS  FOR  CALCIUM. 

Use  separate  portions  of  a  solution  of  a  calcium  salt,  as  CaCl2,  for 
each  of  the  following  tests: 

1.  Add  (NH4)2CO3  T.S.  to  the  calcium  solution.    A  white  preci- 
pitate of  calcium  carbonate,  CaCO3,  will  be  obtained,  which  is  not 
soluble  in  an  excess  of  the  reagent  or  in  alkalis,  but  is  readily  soluble 
in  acids.    Any  soluble  carbonate  will  give  the  same  precipitate. 

2.  Add  NaOH  T.S.  to  the  calcium  solution.    A  white  precipitate 
of  calcium  hydroxide,  Ca(OH)2,  will  be  obtained  if  the  solutions  are 
sufficiently  concentrated.    No  precipitate  is  obtained  from  dilute 
solutions,  as  Ca(OH)2  is  slightly  soluble  in  water.     It  is  less  soluble 


68  THE  METALS 

in  hot  water  than  it  is  in  cold  water.    Ammonium  hydroxide  does 
not  precipitate  calcium  from  solutions  of  calcium  salts. 

3.  Add  H2S04  T.S.  to  the  calcium  solution.    A  white  precipitate 
of  calcium  sulphate,  CaSO4,  will  be  obtained,  unless  the  solution  is 
very  dilute.     CaSO4  is  slightly  soluble  in  water.    Soluble  sulphates, 
as  Na2SO4,  will  give  the  same  precipitate. 

4.  Add  NagHPCX  T.S.  to  the  calcium  solution.    A  white  pre- 
cipitate of  calcium  hydrogen  phosphate,  CaHP04,  will  be  obtained. 
The  precipitate  is  readily  soluble  in  acids. 

5.  Add  (NH4)2C2O4  T.S.  to  the  calcium  solution.    A  white  preci- 
pitate of  calcium  oxalate,  CaC2O4,  will  be  obtained.    The  precipi- 
tate is  readily  soluble  in  mineral  acids  or  in  oxalic  acid,  but  is  not 
soluble  in  acetic  acid. 

6.  Clean  a  piece  of  platinum  wire  by  holding  it  In  a  blue  gas 
flame  until  it  ceases  to  give  any  color  to  the  flame.    Dip  the  wire  in 
the  calcium  solution  and  heat  again.    The  flame  will  be  colored 
yellowish-red. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  CaCl2  +  (NH4)2C03  = 
Ca(C2H302)2  +  Na2C03  = 

2.  Ca(C2H3O2)2  +  NaOH  = 
CaBr2  +  KOH  = 

3.  CaBr2  +  K2SO4  = 
CaCl2  +  (NH4)2S04  = 

4.  CaCl2  +  Na2HPO4  = 
Ca(C2H302)2  +  Na2HP04  = 

5.  CaCl2  +  (NH4)2C2O4  = 
Ca(C2H302)2  +  K2C2O4  = 

STRONTIUM,  Sr11  =  87.63 

Strontium  is  a  yellow  metal  which  readily  oxidizes  in  moist  air 
and  actively  decomposes  water.     It  is  seldom  seen  in  the  metallic 

state. 

IMPORTANT  COMPOUNDS  OF  STRONTIUM. 

Strontium  bromide,  U.  S.  P.,  SrBr26H2O. 
Strontium  carbonate,  SrCO3. 
Strontium  chloride,  SrCl26H2O. 
Strontium  iodide,  U.  S.  P.,  SrI26H2O. 
Strontium  nitrate,  Sr(NO3)2. 

TESTS  FOR  STRONTIUM. 

Use  separate  portions  of  a  solution  of  a  strontium  salt,  as  SrCl2, 
for  each  of  the  following  tests: 


COLLtGt 
of   PHARMACY 

BARIUM  69 

1.  Add  (NH4)2CO3  T.S.  to  the  strontium  solution.    A  white 
precipitate  of  strontium  carbonate,   SrCO3,  will  form,  which  is 
insoluble  in  an  excess  of  the  reagent  or  in  alkalies,  but  is  readily 
soluble  in  acids. 

2.  Add  H2S04  T.S.  to  the  strontium  solution.    A  white  precipi- 
tate of  strontium  sulphate,  SrSO4,  will  form  if  the  solution  is  not  too 
dilute.    The  precipitate  is  insoluble  in  acids  or  alkalies. 

3.  Add  CaSO4  T.S.  to  the  strontium  solution  and  warm  the 
mixture.     A  white  precipitate  of  strontium  sulphate,  SrS04,  will 
form  slowly.    As  CaSO4  is  only  slightly  soluble  in  water,  the  reagent 
is  very  dilute,  so  the  amount  of  the  precipitate  is  small.    As  this 
reagent  will  not  precipitate  calcium  it  can  be  used  to  detect  stron- 
tium in  the  presence  of  calcium. 

4.  Add  Na2HPO4  T.S.  to  the  strontium  solution.    A  white  pre- 
cipitate of  strontium  hydrogen  phosphate,  SrHPO4,  will  form,  which 
is  soluble  in  acids. 

5.  Add  (NH4)2C2O4  T.S.  to  the  strontium   solution.    A  white 
precipitate  of  strontium  oxalate,  SrC2O4,  will  form.    The  precipitate 
is  readily  soluble  in  mineral  acids,  but  only  very  slightly  soluble  in 
acetic  or  oxalic  acid  solutions. 

6.  Hold  a  piece  of  platinum  in  a  non-luminous  gas  flame  until  it 
ceases  to  give  any  color  to  the  flame.    Dip  the  cleaned  wire  in  the 
strontium  solution  and  again  hold  it  in  the  flame.    The  flame  will  be 
colored  deep  red. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  SrCl2  +  (NH4)2C03  = 
Sr(N03)2  +  Na2CO3  = 

2.  Sr(N03)2  +  H2S04  = 
SrCl2  +  Na2SO4  = 

3.  SrCl2  +  CaS04  = 
Sr(N03)2  +  CaSO4  = 

4.  Sr(NO3)2  +  Na2HPO4  = 
SrCl2  +  (NH4)2HPO4  = 

5;    SrCl2  +  (NH4)2C204  = 
Sr(N03)2  +  K2C204  = 

BARIUM,  Bau  =  137.37 

Barium  is  a  yellowish-white  metal  which  is  not  used  in  the  metallic 
state.     Its  compounds  are  poisonous. 

IMPORTANT  COMPOUNDS  OF  BARIUM. 
Barium  carbonate,  BaCO3. 
Barium  chloride,  BaCl22H2O. 
Barium  dioxide,  barium  peroxide,  Ba02. 


70  THE  METALS 

Barium  hydroxide,  Ba(OH)2. 

Barium  nitrate,  Ba(NO3)2. 

Barium  sulphate,  permanent  white,  BaSO4. 

Barium  sulphide,  BaS. 

TESTS  FOR  BARIUM. 

Use  a  separate  portion  of  a  solution  of  a  barium  compound,  as 
BaCl2,  for  each  of  the  following  tests: 

1.  Add  (NH4)2C03  T.S.  to  the  barium  solution.    A  white  pre- 
cipitate of  barium  carbonate,  BaC03,  will  be  obtained.    Other 
soluble  carbonates,  as  Na2CO3,  will  give  the  same  precipitate. 

2.  Add  H2SO4  T.S.  to  the  barium  solution.    A  white  precipitate 
of  barium  sulphate,  BaSO4,  will  be  obtained,  which  is  insoluble  in 
acids  or  alkalies.    Soluble   sulphates,   as  K2SO4,   give  the   same 
precipitate. 

3.  Add  (NH4)2C2O4  T.S.  to  the  barium  solution.    A  white  precip- 
itate of  barium  oxalate,  BaC204,  will  be  obtained.    The  precipitate 
is  soluble  in  acetic  acid  or  in  a  solution  of  oxalic  acid. 

4.  Add  Na2HPO4  T.S.  to  the  barium  solution.    A  white  precipi- 
tate of  barium  hydrogen  phosphate,  BaHPO4,  will  be  obtained. 

5.  Add  K2Cr2O7  T.S.  to  the  barium  solution.    A  yellow  precipi- 
tate of  barium  chromate,  BaCrO4,  will  be  obtained.    This  precipi- 
tate is  soluble  in  nitric  and  hydrochloric  acids,  but  is  insoluble  in 
acetic  acid.    K2Cr04  T.S.  will  give  the  same  precipitate. 

6.  Make  a  flame  test  of  the  barium  solution  with  platinum  wire. 
The  flame  will  be  colored  a  yellowish-green. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  BaCl2  +  (NH4)2C03  = 
Ba(NO3)2  +  Na2C03  = 

2.  Ba(NO3)2  +  H2SO4  = 
BaCl2  +  K2S04  = 

3.  Ba(C2H302)2  +  (NH4)2C204  = 
BaCl2  +  K2C204  = 

4.  BaCl2  +  Na2HP04  = 
Ba(N03)2  +  (NH4)2HP04  = 

5.  BaCl2  +  K2Cr2O7  +  H20  = 
Ba(NO3)2  +  K2Cr04  = 

SAMPLE  NUMBER  10. 

Analysis  of  a  solution  for  group  6  of  the  metals. 

The  solution  may  contain  salts  of  any  of  the  following  metals: 
calcium,  strontium,  barium. 

A .  To  about  10  mils  of  the  sample  add  about  5  mils  of  NH4C1  T.S. 
and  about  5  mils  of  NH4OH  T.S.  and  then  (NH4)2CO3  T.S.  If  no 


ANALYTICAL  TABLES 


71 


precipitate  is  produced,  calcium,  strontium  and  barium  are  absent. 
If  a  precipitate  is  produced,  continue  to  add  (NH^COs  T.S.  until 
precipitation  appears  to  be  complete.  Filter  and  test  the  filtrate. 

B.  Wash  any  precipitate  obtained  in  A  on  the  filter  with  water, 
discarding  the  washings.    Pour  the  same  portion  of  about  5  mils  of 
HC2H3O2  T.S.  through  the  filter  several  times  to  dissolve  the  precipi- 
tate.   Add    a  few   drops  of  K2Cr207  T.S.  to  the  solution.     If  no 
precipitate  is  produced,   barium  is  absent.    A  yellow  precipitate 
shows  the  presence  of  barium.     If  barium  is  present  continue  to 
add  K2Cr2O7  T.S.  until  precipitation  appears  to  be  complete.    Filter 
and  test  the  filtrate.     Divide  the  liquid  or  filtrate  into  two  portions. 

C.  To  one  portion  of  the  liquid  or  filtrate  from  B  add  CaSO4  T.S., 
boil  and  let  stand.     If  no  precipitate   is  produced,  strontium  is 
absent.    A    small    amount    of  white  precipitate,   slowly  formed, 
shows  the  presence  of  strontium. 

D.  To  the  other  portion  of  the  filtrate  from  B,  add  K2S04  T.S., 
boil,  let  stand  for  ten  minutes  and  filter,  to  remove  strontium  if 
present.    To  the  filtrate  add  (NH4)2C2O4  T.S.    If  no  precipitate  is 
produced,  calcium  is  absent.    A  small  amount  of  white  precipitate 
slowly  formed  shows  the  presence  of  calcium. 

SAMPLE  NUMBER  11. 
Analysis  of  a  solution  for  groups  1,  2,  3,  4,  5  and  6  of  the  metals. 

The  solution  should  be  neutral  or  acid  in  reaction  to  litmus.  If  alkaline  add 
HNO3  to  the  portion  to  be  examined  for  metals  until  the  reaction  is  acid.  Add 
HC1  as  long  as  a  precipitate  is  produced  and  filter. 


Precipitate 

Filtrate 

Pb,  Hg', 
Ag. 

Warm  and  pass  H2S  as  long  as  a  precipitate  is  produced.    Filter. 

Examine  for 

Precipitate 

Filtrate 

metals  of 
group  1  by 

Digest  with  (NH4)2Sx 
and  filter. 

Boil  to  expel  H2S.    Add  a  few  drops 
of  HNO3  and  boil  again,  then  add 

I. 

Residue 
Hg",  Pb, 

Filtrate 
Sb,  Sn,  As, 

NH4C1  and  an  excess  of  NH4OH  and 
boil  again.     Filter. 

Bi,  Cu,  Cd. 

Au,  Pt. 

Precipitate  !               Filtrate 

Examine 
for   metals 

Examine 
for   metals 

Al,  Cr,  Fe. 

Examine 

Pass  H2S  as  long  as  a  ppt. 
is  produced.     Filter. 

of  group  2 
by  II. 

of  group  3 
by  III. 

for   metals 
of  group  4 
by  IV. 

Precipitate 
Zn,  Mn, 

To    Ni 

Filtrate 
Boil  to  expel 

V/Uy     Jill* 

BUS,  add  an 

Examine 

excess  of 

for  metals 

(NH4)2C03 

of  group  5 

and  filter. 

byV. 

Precipitate 

Ba,  Sr,  Ca. 

Examine 

for   metals 

of  group  6 
by  VI. 

72 


THE  METALS 


I. 

Examination  of  any  precipitate  produced  by  HOI  in  a  neutral  or 
acid  solution. 


Wash  the  precipitate  on  the  filter  with  cold  water,  discarding  the  washings. 
Pass  a  portion  of  hot  water  through  the  washed  precipitate  on  the  filter. 


Filtrate— PbCl2 

Add  H2SO4  and  cool. 
White  ppt.  PbSO4,  shows 


lead 


Residue— HgCl,  AgCl 

Pass  the  same  portion  of  NH4OH  through  the  residue 
on  the  filter  several  times 


Residue— NH2HgCl+Hg, 

black,  shows 

mercurous  mercury 


Filtrate,  (NH3)3(AgCl)2 
Add  an  excess  of  HNOs, 
white  ppt.,  AgCl,  shows 

silver 


II. 

Examination  of  any  residue  insoluble  in  (NH4)2SX  from  a  precipi- 
tate produced  by  H2S  in  an  acid  liquid,  after  group  1  of  the  metals 
has  been  removed  from  a  solution. 


Wash  the  residue  on  the  filter,  first  with  an  additional  portion  of  (NH4)2SX  and 
then  with  water,  and  drain,  discarding  the  washings. 

Pour  a  portion  of  hot  HNO3  through  the  filter  several  times. 


Residue,  HgS 

Filtrate—  Pb(NO3)2,  Bi(NO3)3,  Cu(NO3)2,  Cd(NO3)2 

Black.  Dissolve  i  Add  cone.  H2SO4  and  boil  until  white  fumes  are  given  off,  dilute 

in  nitro-                                         with  H2O,  cool  and  filter. 

hydrochloric       precipitate,               Filtrate—  Bi2(  SO  4K  CuSO4,  CdSO4 

acid,  boil  to            pVk<an 

expel  Cl,  dilute  ; 

Add  an  excess  of  NH4OH  and  filter. 

with  water,            White, 
filter  and  add    Dissolve  with 

Precipitate, 

Filtrate— 

SnCl2.     A         NH4C2H3O2 

Bi(OH)3 

Cu(NH3)4SO4,  Cd(NH3)4SO4 

white  or  gray         and  add 

White.     Add 

ppt.  HgCl+Hg        KsOiOy. 

NaOH  and 

Divide  into  two  portions. 

shows 

Yellow  ppt., 
PbOO4, 
shows 

SnCl2  on 
filter.  Brown 
or  black 
coloration, 
Bi2O3,    shows 

To  one  portion  !       Pass  H2S. 
add  an  excess       through  the 
of  HC2H3O2     i  other  portion, 
and  then       ;  first  decoloriz- 
K4Fe(CN)6.  A  i  ing  with  KCN 

red  ppt.  or      if  Cu  is  present. 

coloration        A  yellow  ppt., 

Cu2Fe(CN)6,    •      CdS,  shows 

mercuric 

shows 

mercury 

lead 

bismuth 

copper              cadmium 

ANALYTICAL  TABLES 


III. 


73 


Examination  of  the  solution  in  (NH^Sx  of  any  precipitate  pro- 
duced by  H2S  in  an  acid  liquid,  after  group  1  of  the  metals  has  been 
removed  from  a  solution. 


Acidulate  the  (NH4)2SX  solution  with  HC1  and  collect  any  precipitate  formed 
on  a  filter.   Transfer  the  precipitate  to  a  test  tube  and  digest  in  (NH4)2CO3  T.S. 

Filter. 

Filtrate— 

Residue—  Sb2S5,  SnS2,  Au2S3,  PtS2 

(NH4)3AsS4 

Wash  with  water.     Digest 

in  concentrated 

HC1.    Filter. 

Add  excess  of 
HC1  and  pass 

Filtrate—  SbCls,  SnCl2 

Residue  —  Au2S3,  PtS2 

H2S.  A  yellow 

Place  in  H  generator.     Add 

If   light 

colored    Au    and    Pt 

ppt.,  As2S3, 

H2O  and  Zn. 

Conduct  gas 

are  absent.    If  dark  colored,  di- 

shows 

into  AgNOs  for  three  minutes 

gest  in  nitrohydrochloric  acid. 

If  black  ppt. 

Filter    liquid 

Add  H2O  and  filter. 

forms  in 

from  genera- 

To a  part  of 

'Neutralize 

AgNOs  dis- 

tor and  add  few 

the  filtrate  add 

another  part  of 

solve  in  HC1 

drops  to  HgCl2. 

SnCl2  and 

the  filtrate  with 

and  pass  HsS. 

A  white  or 

SnCl4  and  let 

NH4OH,  add 

An  orange 

gray  ppt., 

stand.    A  pur- 

alcohol and 

ppt.,   Sb2§3, 

HgCl4-Hg, 

ple  ppt.,  Au2O, 

NH4C1.    A  yel- 

shows 

shows 

etc.,  shows 

low  ppt., 

(NH4)2PtCl6, 

shows 

arsenic 

antimony 

tin 

gold 

platinum 

rw 

Examination  of  any  precipitate  produced  by  NH4OH,  after  groups 

1,  2  and  3  of  the  metals  have  been  removed  from  a  solution. 

Wash  the  precipitate  on  the  filter  with  water  and  drain,  discarding  the  washings. 
Boil  the  precipitate  with  NaOH  for  several  minutes  and  filter. 

Filtrate, 

Residue  —  Cr(OH)3,  Fe(OH)3 

NaAlO2. 

.   Wash  with  H2O  and  drain, 

discarding  the  washings.    Trans- 

Acidulate 
with  HC1, 

fer  a  part  of  the  residue  to  a  porcelain  crucible  cover,  add  about 
the  same  amount  of  dry  KNOs  and  about  twice  as  much  dry 

then  add  an 

Na2CO3  and  ignite  until  fused.    Cool,  heat  the  mass  with  H2O 

excess  of 

and  filter. 

(NH4)2C03 
and  boil. 

Filtrate, 

Residue  —  Fe2O3 

A  white  ppt., 

xv2CrO4« 

Wash  on  the  filter  with  water,  discarding  the 

A1(OH)3, 

Acidulate 

washings.    Dissolve  in  HC1 

and  add  KCNS. 

shows 

with 

A  red  coloration,  Fe(CNS)3,  shows 

HC2H3O2 

and  add 

iron 

Pb(C2H302)2 

If  iron  is  present  test  separate  portions  of  the 

A  yellow 

original  solution  as  follows  : 

ppt  .  ,   PbCrO 

*'    To  one  portion  add 

To    another    portion 

shows 

HC1  and  K3Fe(CN)6. 

add  HC1  and  KCNS. 

A  dark  blue  ppt., 

A  red  coloration, 

Fe3(Fe(CN)< 

)2,  shows 

Fe(CNS)3,  shows 

aluminum 

chromium 

ferrous 

iron 

ferric  iron 

74 


THE  METALS 
V. 


Examination  of  any  precipitate  produced  by  H2S  in  an  alkaline 
liquid,  after  groups  1,  2,  3  and  4  of  the  metals  have  been  removed 
from  a  solution. 

Wash  the  precipitate  on  the  filter  with  water  and  drain,  discarding  the  wash- 
ings. Pass  the  same  portion  of  HC1  through  the  filter  several  times. 


Filtrate— ZnCl2,  MnCl2. 


Residue — CoS,  NiS. 


Boil  to  expel  H2S,  cool,  add 

If  light  colored,  cobalt  and  nickel  are  absent. 

excess  of  NaOH  and  filter. 

If   dark   colored,    dissolve   in   nitro-hydro- 

Filtrate, 
NftaZnOj, 

Pass  H2S. 

Residue, 
Mn(OH)2. 

Boil  with  cone. 

chloric  acid,  add  an  excess  of  NaOH,  collect 
the  ppt.  produced  on  a  filter  and  wash,  dis- 
carding the  washings.     Dissolve  the  ppt.  in 
HC2H3O2,  add  an  excess  of  KC1  and  of  KNO2, 

A   white   ppt. 

HNO3  and 

let  stand  for  half  an  hour  and  filter. 

ZnS,  shows 

PbO2   and   let 
stand.    A  pink 
supernatant 

Precipitate, 

K3Co(N02)6 

Filtrate,  Ni(C2H3O2)2 
Add  an  excess  of 

liquid  shows 

yellow,  shows 

NaOH.    Light  green 

ppt.  Ni(OH)2,  shows 

zinc               manganese 

cobalt 

nickel 

VI. 

Examination  of  any  precipitate  produced  by  (NH4)2C03,  after 
groups  1,  2,  3,  4  and  5  of  the  metals  have  been  removed  from  a 
solution. 

Wash  the  precipitate  on  the  filter  with  water  and  drain,  discarding  the 
washings. 

Dissolve  the  washed  precipitate  in  HC2H3O2,  add  an  excess  of  K2Cr2O:  and 
filter. 


Precipitate, 
BaCrO; 

Yellow 
shows 


barium 


Filtrate— Sr(C2H3O 2) 2,  Ca(C2H3O2)2. 

Divide  into  two  portions. 
To  one  portion  add         To  the  other  portion,  add  K2SO4, 


CaSO4,  boil  and  let 

stand.    White  ppt. 

SrSO4,  shows 

strontium 


boil,  cool  and  filter,  if  necessary. 
To  the  filtrate  add  (NH4)2C2O4. 
White  ppt.  CaC2O4,  shows 

calcium 


GROUP   7. 
Metals  not  precipitated  by  any  group  reagent: 

Magnesium,    Mg;    potassium,    K;    sodium,    Na;    lithium,    Li; 
ammonium,  (NEU)1. 

MAGNESIUM,  Mgu  =  24.32 

Magnesium  is  a  silvery-white  metal  of  low  specific  gravity  which  is 
permanent  in  dry  air  but  oxidizes  in  moist  air.     It  burns  in  the  air, 


MAGNESIUM  75 

giving  an  intense  white  light.      It  is  a  constituent  of  flashlight 
powders,  but  is  most  important  in  its  compounds. 

IMPORTANT  COMPOUNDS  OF  MAGNESIUM. 

Magnesium  carbonate,  native  as  magnesite,  MgCO3. 

Magnesium  sub-carbonate,  magnesium  carbonate,  U.  S.  P.,  magnesia 

alba,  (MgCO3)4Mg(OH)25H2O. 
Magnesium  chloride,  N.F.,  MgCl26H2O. 
Magnesium  hydroxide,  Mg(OH)2. 

Magnesium  oxide,  U.  S.  P.,  magnesia,  calcined  magnesia,  MgO. 
Heavy  magnesium  oxide,  U.  S.  P.,  heavy  magnesia,  MgO. 
Magnesium  sulphate,  U.  S.  P.,  epsom  salt,  MgSO47H2O. 
Magnesium   silicates,   talcum,    asbestos,   meerschaum,   soapstone; 

variable  in  composition. 

TESTS  FOR  MAGNESIUM. 

Use  a  separate  portion  of  a  solution  of  a  salt  of  magnesium,  as 
MgCl2,  for  each  of  the  following  tests: 

1.  Add  (NH4)2CO3  T.S.  to  the  magnesium  solution.    A  white 
precipitate  of  magnesium-ammonium  carbonate,  MgCO3(NH4)2CO3, 
is  formed  when  the  reagent  is  in  excess.    No  precipitate  will  form  if 
excessive  amounts  of  ammonium  salts  are  present. 

To  a  portion  of  the  magnesium  solution  add  an  equal  volume  of 
NH4C1  T.S.  and  then  (NH4)2CO3  T.S.  Soluble  magnesium-ammo- 
nium chloride  is  formed  and  this  is  not  precipitated  by  the  carbonate. 

2.  Add  Na2C03  T.S.  to  the  magnesium  solution.    A  white  pre- 
cipitate of  magnesium  sub-carbonate,  principally  (MgCO3)4Mg(OH)2 
is  formed.    The  precipitate  is  soluble  in  solutions  of  ammonium 
salts  and  will  not  form  if  they  are  present. 

3.  Add  NH4OH  T.S.  to  the  magnesium  solution.    Half  of  the 
magnesium   is  precipitated   as  the  white   gelatinous  magnesium 
hydroxide,   Mg(OH)2.    The   other  half  forms   a  double   salt,   as 
(NH4)2MgCl4,  and  remains  in  solution.   The  precipitate  forms  soluble 
double  compounds  with   ammonium   salts   and   no   precipitation 
occurs  if  an  excess  of  ammonium  salts  is  present  in  the  solution. 

4.  Add  NaOH  T.S.  to  the  magnesium  solution.    A  gelatinous 
white  precipitate  of  magnesium  hydroxide,  Mg(OH)2,  will  form. 
The  precipitate  is  soluble  in  ammonium  salts  and  will  not  form  if 
they  are  present. 

5.  To  a  portion  of  the  magnesium  solution  add  NH4C1  T.S., 
NH4OH  T.S.,  and  Na2HPO4  T.S.  in  the  order  given.    A  white  crys- 
talline precipitate  of  magnesium-ammonium  phosphate,  MgNH4PO4, 
will  form  slowly.    NH4C1  does  not  take  part  in  the  reaction,  but  is 
used  to  prevent  precipitation  of  magnesium  hydroxide  by  NH4OH. 


76  THE  METALS 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  MgCl2  +  (NH4)2C03  = 
MgS04  +  (NH4)2C03  - 
MgCl2  +  NH4C1  = 

2.  MgCl2  +  K2C03  = 
MgSO4  +  Na2CO3  = 

3.  MgBr2  +  NH4OH  = 
MgS04  +  NH4OH  = 

4.  MgSO4  +  NaOH  = 
MgCl2  +  KOH  = 

5.  MgCl2  +  NH4OH  +  Na2HPO4  = 
MgSO4  +  NH4OH  +  (NH4)2HP04  = 

POTASSIUM  (KALIUM),  Ks  =  39.10 

Potassium  is  the  most  electropositive  of  the  metals  which  are  found 
in  any  considerable  quantity.  It  is  a  soft  white  metal  which  must  be 
kept  under  kerosene,  as  it  rapidly  oxidizes  and  takes  fire  in  contact 
with  water  or  when  exposed  to  the  air.  The  element  is  but  little 
used  in  the  metallic  state.  The  salts  of  potassium  are  very  impor- 
tant and  all  of  them  are  more  or  less  soluble  in  water. 

IMPORTANT  COMPOUNDS  OF  POTASSIUM. 

Potassium  acetate,  U.  S.  P.,  KC2H3O2. 

Potassium  bicarbonate,  U.  S.  P.,  saleratus,  KHCO3. 

Potassium  bitartrate,  U.  S.  P.,  cream  of  tartar,  KHC4H4Oe. 

Potassium  bromate,  KBrO3. 

Potassium  bromide,  U.  S.  P.,  KBr. 

Potassium  carbonate,  U.  S.  P.,  potash,  pearlash,  salt  of  tartar,  K2CO3. 

Potassium  chlorate,  U.  S.  P.,  KC1O3. 

Potassium  chloride,  N.F.,  KC1. 

Potassium  chromate,  K2CrO4. 

Potassium  cyanide,  prussiate  of  potash,  KCN. 

Potassium  dichromate,  K2Cr2O?. 

Potassium  ferricyanide,  red  prussiate  of  potash,  K3Fe(CN)6. 

Potassium  ferrocyanide,  yellow  prussiate  of  potash,  K4Fe(CN)6. 

Potassium  hydroxide,  U.  S.  P.,  potassium  hydrate,  potassa,  caustic 

potash,  KOH. 

Potassium  hypophosphite,  U.  S.  P.,  KPH2O2. 
Potassium  iodide,  U.  S.  P.,  KI. 
Potassium  nitrate,  U.  S.  P.,  saltpetre,  nitre,  KNO3. 
Potassium  perchlorate,  KC1O4. 
Potassium  permanganate,  U.  S.  P.,  KMn04. 
Potassium  sulphate,  N.F.,  K2SO4. 


SODIUM  77 

Potassium  sulphocyanate,  potassium  thiocyanate,  KCNS. 
Potassium  and  sodium  tartrate,  U.  S.  P.,  rochelle  salt,  KNaCJ^Oe 

4H20. 
Sulphuretted  potash,  U.S.P.,  liver  of  sulphur,  composition  indefinite. 

TESTS  FOR  POTASSIUM. 

Use  a  separate  portion  of  a  solution  of  a  potassium  salt,  as  KC1 
for  each  of  the  following  tests: 

1.  Add  H2C4H406  T.S.  to  the  potassium  solution.    A  white  pre- 
cipitate of  potassium  bi-tartrate,  KHC4H4O6,  will  form  if  the  solu- 
tion is  not  too  dilute.     The  addition  of  an  equal  volume  of  alcohol 
increases  the  delicacy  of  the  test,  as  the  precipitate  is  less  soluble  in 
diluted  alcohol  than  it  is  in  water.     The  precipitate  is  soluble  in 
alkalies,  and  will  not  form  if  they  are  present  in  the  solution  until  an 
excess  of  the  reagent  has  been  added.     Sodium  hydrogen  tartrate, 
NaHC^Oe,  will  produce  the  same  precipitate  in  neutral  or  acid 
solutions. 

2.  Acidulate  a  portion  of  the  potassium  solution  with  HC2H3O2 
T.S.,  add  Na3Co(NO2)6  T.S.  and  let  stand.    A  yellow  precipitate  of 
potassium-cobaltic  nitrite,  K3Co(NO2)e,  will  form    slowly,  if  the 
solution  is  not  too  dilute. 

3.  Heat  some  of  the  potassium  solution  on  platinum  wire  in  a 
blue    gas  flame  and  observe  the  color  of  the  flame,  directly  and 
through  blue  glass.    A  characteristic  violet  flame,  visible  through 
blue  glass,  is  obtained. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  KC1  +  H2C4H406  = 
K2S04  +  NaHC4H4O6  = 

2.  KC2H302  +  Na3Co(NO2)6  = 
KBr  +  Na3Co(N02)6  = 

SODIUM  (NATRIUM),  Na1   -   23.0 

Sodium  is  a  soft  white  metal,  closely  resembling  potassium. 
It  decomposes  water  and  rapidly  oxidizes  in  the  air,  so  it  must  be 
kept  under  kerosene  or  some  other  liquid  which  does  not  contain 
oxygen.  The  compounds  of  sodium  are  very  important  and  all  of 
them  are  soluble  in  water. 

IMPORTANT  COMPOUNDS  OF  SODIUM. 

Sodium  acetate,  U.  S.  P.,  NaC2H3023H2O. 
Sodium  arsenate,  U.  S.  P.,  Na2HAs047H2O. 
Exsiccated  sodium  arsenate,  U.  S.  P.,  Na2HAsO4. 


78  THE  METALS 

Sodium  bicarbonate,  U.  S.  P.,  baking  soda,  NaHCO3. 

Sodium  bisulphite,  NaHSO3. 

Sodium  borate,  U.  S.  P.,  sodium  tetraborate,  sodium  pyroborate, 

borax,  Na^CMOHA 
Sodium  bromide,  U.  S.  P.,  NaBr. 

Sodium  carbonate,  sal  soda,  washing  soda,  Na2CO310H20. 
Monohydrated  sodium  carbonate,  U.  S.  P.,  Na2CO3H2O. 
Sodium  chlorate,  NaC103. 
Sodium  chloride,  U.  S.  P.,  salt,  NaCl. 
Sodium  cyanide,  U.  S.  P.,  NaCN. 
Sodium  hydroxide,  U.  S.  P.,  sodium  hydrate,  caustic  soda,  soda, 

NaOH. 

Sodium  hypochlorite,  NaOCl. 
Sodium  hypophosphite,  U.  S.  P.,  NaPH2O2H2O. 
Sodium  iodide,  U.  S.  P.,  Nal. 
Sodium  nitrate,  U.  S.  P.,  Chili  saltpetre,  NaNO3. 
Sodium  nitrite,  U.  S.  P.,  NaNO2. 
Sodium  oxalate,  Na2C2O4. 
Sodium  perborate,  U.  S.  P.,  NaBO34H2O. 
Sodium  peroxide,  Na2O2. 
Tri-sodium  phosphate,  Na3P04. 
Sodium   phosphate,    U.   S.    P.,   di-sodium    hydrogen    phosphate, 

Na,HP0412H2O. 

Exsiccated  sodium  phosphate,  U.  S.  P.,  Na2HPO4. 
Sodium  di-hydrogen  phosphate,  NaH2PO4. 
Sodium  pyrophosphate,  Na4P2O7lOH2O. 

Sodium  silicates,  etc.,  water  glass,  soluble  glass,  Na2Si03,  Na2Si409. 
Sodium  sulphate,  U.  S.  P.,  Glauber's  salt,  Na2SO410H20. 
Sodium  sulphide,  Na2S9H2O. 
Sodium  sulphite,  Na2SO37H2O. 
Exsiccated  sodium  sulphite,  U.  S.  P.,  Na2S03. 
Sodium  tartrate,  Na2C4H4062H2O. 
Sodium  thiosulphate,  U.  S.  P.,  "hyposulphite  of  soda,"  Na^SASHA 

TEST  FOR  SODIUM. 

As  all  compounds  of  sodium  are  soluble  in  water  to  an  appreciable 
extent,  there  are  no  dependable  precipitation  tests  for  the  metal. 
The  flame  test  is  very  sensitive,  however,  if  properly  applied. 

Heat  a  piece  of  platinum  wire  in  the  blue  gas  flame  until  it  ceases 
to  impart  any  color  to  the  flame.  Dip  the  wire  in  a  solution  of 
sodium  chloride  and  heat  again,  observing  the  color  directly  and 
through  blue  glass.  An  intense  yellow  color  is  imparted  to  the 
flame,  but  this  color  is  cut  off  by  blue  glass. 

As  sodium  is  widely  distributed  in  small  amounts,  it  is  necessary  to 


AMMONIUM  79 

be  especially  careful  in  cleaning  the  platinum  wire  before  applying 
the  flame  test  for  it.  Also  it  is  necessary  to  test  the  original  solution 
for  sodium,  in  separations,  because  some  of  the  reagents  used  con- 
tain its  salts.  The  flame  test  for  sodium  is  so  delicate  that  the  metal 
is  often  found  and  reported,  when  its  compounds  are  present  only 
as  impurities  in  a  substance. 

LITHIUM,  Li1  =  6.94 

Lithium  is  a  rare  metal  which  is  only  of  importance  in  a  few  of  its 
salts,  used  in  medicine.  Its  compounds  are  generally  colorless  and 
soluble  in  water. 

IMPORTANT  COMPOUNDS  OF  LITHIUM. 

Lithium  bromide,  U.  S.  P.,  LiBr. 
Lithium  carbonate,  U.  S.  P.,  Li2C03. 
Lithium  chloride,  LiCl. 

TESTS  FOR  LITHIUM. 

Use  a  separate  portion  of  a  solution  of  a  lithium  salt,  as  LiCl,  for 
each  of  the  following  tests: 

1.  Add  Na2CO3  T.S.  to  the  lithium  solution.    A  white  precipi- 
tate of  lithium  carbonate,  Li2CO3,  is  formed  if  the  solution  is  suffi- 
ciently concentrated. 

2.  Add  Na2HPO4  T.S.  and  NaOH  T.S.  to  the  lithium  solution 
and  boil  the  mixture.    A  white  precipitate  of  lithium  phosphate, 
Li3P04,  will  form,  if  the  solution  is  not  too  dilute. 

3.  Heat  some  of  the  lithium  solution  on  platinum  wire  in  a  blue 
gas  flame  and  observe  the  color  of  the  flame,  directly  and  through 
blue  glass.    The  flame  is  colored  a  deep  red,  which  is  visible  through 
a  single  piece  of  blue  glass  of  ordinary  thickness,  but  is  not  visible 
through  several  thicknesses  of  blue  glass. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  LiCl  +  Na2CO3  = 
LiBr  +  K2C03  = 

2.  LiBr  +  Na2HPO4  = 
LiCl  +  (NH4)2HP04  = 

AMMONIUM,  (NH4)1  =  18.042 

Ammonium  is  a  positive  compound  radical  which  forms  a  series  of 
salts  that  are  very  similar  to  the  salts  of  potassium  and  sodium. 
All  ammonium  salts  are  volatilized  or  decomposed  by  heat  and  most 
of  them  are  soluble  in  water. 


80  THE  METALS 

IMPORTANT  AMMONIUM  COMPOUNDS. 

Ammonia,  NH3. 

Ammonium  acetate,  NH4C2H3O2. 

Ammonium  bromide,  U.  S.  P.,  NELiBr. 

Ammonium  carbonate,  U.S.P.,  sal  volatile,  NH4HCO3+NH4NH2CO2. 

Ammonium  chloride,  U.  S.  P.,  sal  ammoniac,  NH4C1. 

Ammonium  hydroxide,  NH4OH  (in  solution  only). 

Ammonium  hypophosphite,  N.  F.,  NH4PH2O2. 

Ammonium  iodide,  U.  S.  P.,  NH4I. 

Ammonium  nitrate,  NH4NO3. 

Ammonium  oxalate,  (NH4)2C2O4H2O. 

Ammonium  phosphate,  N.  F.,  (NH4)2HP04  +  NH4H2PO4. 

Ammonium  sulphate,  (NH4)2SO4. 

Ammonium  sulphide,  (NH4)2S. 

Ammonium  poly  sulphide,  yellow  ammonium  sulphide,  (NH4)2SX. 

TESTS  FOR  AMMONIUM. 

Use  a  separate  portion  of  a  solution  of  an  ammonium  salt,  as 
NH4C1,  for  each  of  the  following  tests: 

1.  Add  NaOH  T.S.  to  the  ammonium  solution  until  the  reaction 
is  alkaline  to  litmus,  and  boil  the  mixture.    Ammonia  gas,  NH3,  is 
evolved.    Observe  the  odor  of  ammonia  in  the  vapor.    Moisten  a 
piece  of  red  litmus  paper  and  hold  it  in  the  vapor:  the  color  will 
be  changed  to  bluel    Dip  a  glass  rod  in  concentrated  HC1  and  hold 
it  in  the  vapor:  white  fumes  of  ammonium  chloride,  NH4C1,  are 
formed. 

2.  Add  H2C4H4O6  T.S.  to  the  ammonium  solution.    A  white  pre- 
cipitate of  ammonium  bi-tartrate,  NH4HC4H4O6,  will  be  formed,  if 
the  solution  is  not  too  dilute.    The  precipitate  is  soluble  in  alkalies. 

3.  Acidulate  a  portion  of  the  ammonium  solution  with  HC2H3O2 
T.S.,  add  Na3Co(NO2)6T.S.  and  let  stand.    A  yellow  precipitate  of 
ammonium-cobaltic  nitrite,  (NH^aCoCNO^e,  will  form  slowly. 

4.  Add  alkaline  HgK2I4  T.S.  (Nessler's  reagent)  to  the  ammo- 
nium solution.     A  brown  precipitate  of  nitrogen  dimercuric  iodide, 
NHgJ,  is  formed  if  the  solution  is  concentrated.    With  dilute  solu- 
tions of  ammonium  salts  this  reagent  gives  a  deep  brown  color,  but 
no  precipitate.    Several  of  the  metals  interfere  with  this  reaction. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  NH4C1  +  NaOH  = 
(NH4)2S04  +  KOH  = 
NH3  +  HC1  = 

2.  NH4Br  +  H2C4H4O6  = 
NH4N03  +  NaHC4H4O6  = 


AMMONIUM  81 

3.  NH4C1  +  Na3Co(N02)6  - 
NH4N03  +  Na3Co(NO2)6  = 

4.  NH4C1  +  2HgK2I4  +  4KOH  =  NHg2I  +  7KI  +  KC1  +  4H2O 

SAMPLE  NUMBER  12. 
Analysis  of  a  solution  for  group  7  of  the  metals. 

The  sample  may  contain  compounds  of  any  of  the  following 
metals:    magnesium,  lithium,  sodium,  potassium,  ammonium. 

A.  Acidulate  about  5  mils  of  the  sample  with  HC1  T.S.,  evapo- 
rate the  liquid  to  a  small  volume  on  a  water-bath,  transfer  to  a  porce- 
lain crucible  and  ignite  gently,  to  volatilize  any  ammonium  com- 
pounds.    Dissolve  the  residue  in  about  10  mils  of  water  in  a  beaker, 
adding  a  few  drops  of  HC1  T.S.,  to  obtain  a  clear  solution,  if  neces- 
sary. 

B.  Add  an  excess  of  NaOH  T.S.  to  the  solution  obtained  in  A 
and  boil.    If  no  precipitate  is  produced,  magnesium  is  absent. 
A  white  precipitate,  Mg(OH)2,  shows  the  presence  of  magnesium. 
Filter  and  test  the  filtrate. 

C.  To  the  clear  filtrate  obtained  in  J5,  add  Na2HPO4  T.S.  and 
boil.     If  no  precipitate  is  produced,  lithium  is  absent.     A  white 
precipitate,  Li3PO4,  which  may  appear  as  no  more  than  a  turbidity 
in  the  liquid,  shows  the  presence  of  lithjki. 

D.  Hold  the  end  of  a  piece  of.  platfflfm  wir^P  the  blue  flame 
of  a  gas  burner  until  the  flame  fro  longdfciows  a  color  from  the  wire. 
Dip  the  wire  in  the  original  sojpion  anBiold  it  yi  the  flame  again. 

If  the  flame  is  bright  yello«  when  Bked  at  directly  but  shows 
no  color  through  blue  glassRdium  Wpresent  and  potassium  is 


absent. 

^p 
thick  blue  glass,  potassium  i^p-esent  and  sodium  is  absent. 


JP 


If  the  flame  is  violet  whenkat  cmrectly  and  through  double 


If  the  flame  is  brightMellow  when  looked  at  directly  and  shows 
a  violet  color  through  crouble  thick  blue  glass,  both  sodium  and 
potassium  are  present. 

If  lithium  is  present  these  flame  reactions  are  affected  by  its 
deep  red  flame,  but  this  can  be  allowed  for  after  experience  with  the 
tests  described  above. 

E.  To  a  portion  of  the  original  solution  add  an  excess  of  NaOH 
T.S.  and  boil.  If  no  odor  of  ammonia  can  be  detected  and  the  steam 
will  not  turn  red  litmus  paper  blue,  ammonium  is  absent.  If  the 
steam  smells  of  ammonia,  or  turns  red  litmus  paper  blue,  ammonium 
is  present. 


82 


THE  METALS 


SAMPLE  NUMBER  13. 
Analysis  of  a  solution  for  all  groups  of  the  metals. 

Separation  of  the  metals  into  groups. 


The  solution  should  be  neutral  or  acid  in  reaction  to  litmus.  If  alkaline  add 
HNOs  to  the  portion  to  be  examined  for  metals  until  the  reaction  is  acid.  Add 
HC1  as  long  as  a  precipitate  is  produced  and  filter. 


Precipitate 

Pb,  Hg', 

Ag. 

Examine 

for  metals 

of  group  1 


Filtrate 
Warm  and  pass  H2S  as  long  as  a  precipitate  is  produced.     Filter. 


Precipitate 

Digest  with 
(NH4)2SX  and  filter. 


Residue 

Hg",  Pb, 

Bi,Cu,Cd. 

Examine 

for  metals 

of  group 

2  by  II. 


Filtrate 

Sb,Sn,As, 

Au,  Pt. 

Examine 

for  metals 

of  group 

3  by  III. 


Filtrate 

Boil  to  expel  H2S .  Add  a  few  drops  of  HNO3 
and  boil  again,  then  add  NH4C1  and  an  excess 
of  NH4OH  and  boil  again.  Filter. 

Precipitate  Filtrate 

Al,  Cr,  Fe. 


Examine 

for  metals 

of  group  4 

by  IV. 


Pass  H2S  as  long  as  a  ppt.  is  pro- 
duced.    Filter. 


Precipitate 
Zn,  Mn, 
Co,Ni. 

Examine 
for  metals  ,piitate 

byT  5B».Sr,Ca. 

Examine 

for  metals 

of  group  6 

by  VI. 


Filtrate 

Boil  to  expel  H2S, 

add  an  excess  of 
(NH4)2CO3  and  filter. 

Filtrate 
Mg,  Li, 
Na,  K, 

(NH4)'. 

Examine 

for  metals 

of  group  7 

by  VII. 


Examination  of  any  precipitate  produced  by  HC1  in  a  neutral  or 
acid  solution. 


Wash  the  precipitate  on  the  filter  with  cold  water,  discarding  the  washings. 
Pass  a  portion  of  hot  water  through  the  washed  precipitate  on  the  filter. 


Filtrate— PbCl2 

Add  H2SO4  and  cool. 
White  ppt.  PbSO4,  shows 


lead 


Residue— HgCl,  AgCl 

Pass  the  same  portion  of  NH4OH  through  the  residue 
on  the  filter  several  times 


Residue— NH2HgCl-f-Hg, 

black,  shows 

mercurous  mercury 


Filtrate,  (NH3) 3 (AgCl)  2 

Add  an  excess  of  HNO3, 
white  ppt.,  AgCl,  shows 

silver 


ANALYTICAL  TABLES 


83 


II. 

Examination  of  any  residue  insoluble  in  (NH^Sx  from  a  precipi- 
tate produced  by  H2S  in  an  acid  liquid,  after  group  1  of  the  metals 
has  been  removed  from  a  solution. 

Wash  the  residue  on  the  filter,  first  with  an  additional  portion  of  (NH4)2Sx  and 
then  with  water,  and  drain,  discarding  the  washings. 

Pour  a  portion  of  hot  HNOs  through  the  filter  several  times. 


-Bi2(SO4)3,  CuSO4,  CdSO4 
:cess  of  NH4OH  and  filter. 

Filtrate— 
Cu(NH3)4SO4,  Cd(NH3)4SO4 

Divide  into  two  portions. 

To  one  portion        Pass  H2S 
add  an  excess      through  the 
of  HC2H3O2 
and  then 

K4Fe(CN)6. 

A  red  ppt.  or 

coloration, 


•Residue,  HgS 

Filtrate—  Pb(N03)2,  Bitt 

Clack.  Dissolve 

Add  cone.  H2SO4  and  boil  until 

in  nitro- 

with  H2O,  c 

hydrochloric 
acid,   boil  to 
expel  Cl,  dilute 

Precipitate, 
PbS04 

Filtrate- 
Add  an  e: 

with  water, 
filter  and  add 
SnCl2.     A 

White. 
Dissolve  with 
NH4C2H3O2 

Precipitate, 
Bi(OH)3 

white  or  gray 

and  add 

White.     Add 

ppt.HgCl-fHg 

K2Cr2O7. 

NaOH  and 

shows 

Yellow  ppt., 
PbCrO4, 

SnCl2  on  filter. 
Brown  or 

shows 

black  colora- 

tion, Bi2O3, 

• 

shows 

mercuric 

J 

mercury 

lead 

bismuth  3s 

other  portion, 
first  decoloriz- 
ing with  KCN 
if  Cu  is  present. 
A  yellow  ppt., 
CdS,  shows 


Examination  of  the  solution  in  (NH4)2SX  <> 
duced  by  H2S  in  an  acid  liquid,  after  group  1  < 
removed  from  a  solution, 


precipitate  pro- 
metals  has  been 


Acidulate  the  (NH4)2SX  ^lution  with  HC1  and  collect  an 
on  a  filter.  Transfer  the  precipitate  1<>  a  test  Ujjfc  andrnges 
Filter. 


y  precipitate  formed 
gestin(NH4)2CO3T.S. 


Filtrate— 

»  Residue—  Sb^Ss,  SnS  ,,  Au,>S3,  PtS2 

(NH4)3AsS4 

water.     Digest  in  concentrated  HC1.     Filter. 

Add  excess  of 
HC1  and  pass 

Filtrate—  SbCl3,  SnCl2 

Residue  —  Au2S3,  PtS2 

H2S.  A  yellow 

Place  in  H  generator.     Add 

If    light    colored    Au    and    Pt 

ppt.,  As2S3, 

H2O  and  Zn.     Conduct  gas 

are  absent.    If  dark  colored,  di- 

shows 

into  AgNOs  for  three  minutes 

gest  in  nitrohydro  chloric  acid. 

If  black  ppt. 

Filter    liquid 

Add  H2O  and  filter. 

forms  in 

from  genera- 

To a  part  of 

Neutralize 

AgNO3  dis- 

tor and  add  few 

the  filtrate  add 

another  part  of 

solve  in  HC1 

drops  to  HgCl2. 

SnCl2  and 

the  filtrate  with 

and  pass  H^S. 
An  orange 
ppt.,  Sb2S3, 

A  white  or 
gray  ppt., 
HgCH-Hg, 

SnCl4  and  let 
stand.    A  pur- 
ple ppt.,  Au2O, 

NH4OH,  add 
alcohol  and 
NH4C1.    A  yel- 

shows 

shows 

etc.,  shows 

low  ppt., 

(NH4)2PtClfl, 

shows 

arsenic 

antimony 

tin 

gold                 platinum 

84 


THE  METALS 


IV. 

Examination  of  any  precipitate  produced  by  NH4OH,  after  groups 
1,  2  and  3  of  the  metals  have  been  removed  from  a  solution. 


Wash  the  precipitate  on  the  filter  with  water  and  drain,discarding  the  washings. 
Boil  the  precipitate  with  NaOH  for  several  minutes  and  filter. 


Filtrate, 
NaAlO2. 

Acidulate 
with  HC1, 

then  add  an 

excess  of 

(NH4)2CO3 

and  boil. 

Awhiteppt., 

Al(OH),, 

shows 


Residue— Cr (OH) 3,  Fe(OH)3  % 

Wash  with  H2O  and  drain,  discarding  the  washings.    TransA 

iF 

Cool,  heat  the  mass  with  H2O 


f er  a  part  of  the  residue  to  a  porcelain  crucible  cover,  add  abou 
the  same  amount  of  dry  KNO3  and  about  twice  as  much  drv 


Na2CO3  and  ignite  until  fused, 
and  filter. 


aluminum 


Filtrate, 
K2CrO4. 

Acidulate 

with 

HC2H3O2 

and  add 

Pb(C2H3O2)2. 

A  yellow 

ppt.,   PbCrO4, 

shows 


Residue — Fe2O3 

Wash  on  the  filter  with  water,  discarding  the 
washings.  Dissolve  in  HC1  and  add  KCNS. 
A  red  coloration,  Fe(CNS)i,  shows 

iron 

If  iron  is  present  test  separate  portions  of  the 
original  solution  as  follows: 

To   another 


To  one  portion  add 

HC1  and  K3Fe(CN)6. 

A  dark  blue  ppt., 

(Fe(CN)6)2,  shows 

ferrous  iron 


Examination 
liquid,  after  groups 
from  a  solution. 


portion 
add  HC1  and  KCNS. 
A  red  coloration, 
Fe(CNS)3,  shows 

ferric  iron 


Wash  the  precipitate  on  the 
ings.    Pass  the  same  portion  of 

Filtrate— ZnCl2>  MnCl2. 


by  H2S  in  an  alkaline 
have  been  removed 


,  discarding  the  wash- 
veral  times. 

CoS,  NiS. 


Boil  to  expel  H2S,  cool,  add 

If  light  colored,  robalt  and  nickel  are  absent. 

excess  of  NaOH  and  filter. 

If   dark    colored,    dissolve    in   nitrohydro- 

Filtrate, 
NfeZnOi. 

Residue, 
Mn(OH)2. 

chloric  acid,  add  an  excess  of  NaOH,  collect 
the  ppt.  produced  on  a  filter  and  wash,  dis- 
carding the  washings.     Dissolve'  the  ppt.  in 

Pass  H2S. 

Boil  with 

HC2H3O2,  add  an  excess  of  KC1  and  of  KNO2, 

A  white  ppt. 

HNO3  and 

let  stand  for  half  an  hour  and  filter. 

ZnS,  shows 

PbO2  and  let 
stand.    A  pink 
supernatant 

Precipitate, 

K3Co(NO2)6, 

Filtrate,  Ni(C2H3O2)2 
Add  an  excess  of 

liquid  shows 

yellow,  shows 

NaOH.     Light  green 

ppt.  Ni(OH)2,  shows 

zinc 

manganese 

cobalt 

nick*l 

ANALYTICAL  TABLES 
VI. 


85 


Examination  of  any  precipitate  produced  by  (NH4)2CO3,  after 
groups  1,  2,  3,  4  and  5  of  the  metals  have  been  removed  from  a 
solution. 


Wash  the  precipitate  on  the  filter  with  water  and  drain,  discarding  the  wash- 
incs.  Dissolve  the  washed  precipitate  in  HC2H3O2,  add  an  excess  of  K2Cr2O? 
arm  filter. 


Precipitate, 
BaCrO4 

Yellow 
shows 


barium 


Filtrate— Sr(C2H3O2)2,  Ca(C2H3O2)2. 
Divide  into  two  portions. 

To  one  portion  add 

CaSO4,  boil  and  let 

stand.     White  ppt. 

SrSO4,  shows 


strontium 


To  the  other  portion,  add  K2SO4, 
boil,  cool  and  filter,  if  necessary. 
To  the  filtrate  add  (NH4)2C2O4. 
White  ppt.  CaC2O4,  shows 


calcium 


VII. 

Examination   of   a   solution  for  metals   not  precipitated   with 
groups  1,  2,  3,  4,  5  or  6. 


rate  the  liquid  to 
Add  an  excess  of 


Add  an  excess  of  HC1  to  the  filtrat 
dryness  and  ignite.    Dissolve  the 
NaOH  and  boil.    A  white  ppt., 


Filter  out  any  Mg(OH)2,  add  NaM  D4  and  b<  o*ppt.,  Li3PO4,  shows*' 

lithium.  / 

Make  a  flame  test  of  the  oi^Bial  solution  on  platinum  \\  ire. 

A  yellow  flame,  not  visibj^^rouH  ^pme,    visible    through 

ck  blue  glass  shows 

potassium 

of  the  ol^nal  solution  and  boil.    An 
ammonia,  shows 

w 

ammonium 


blue  glass  show 
sodium  J 

Add  an  excess  of  NaOJ[ 
alkaline  vapor,  NH3,  havm 


THE  ACIDS. 


The  acid  radicals  cannot  be  separated  from  each  other  by  grouping  \ 
as  is  done  with  the  metals,  and  they  must  often  be  tested  for  sepa- 
rately in  the  original  sample.    The  reactions  of  the  soluble  salts  of 
the  most  important  acids  are  given  here.    Some  others  will  be 
noticed  later  when  the  preliminary  treatment  of  solids  is  given. 

Three  groups  of  the  acids  are  recognized,  as  follows : 

Group  A. — Acids  whose  radicals  are  precipitated  by  silver  nitrate, 
in  the  presence  of  nitric  acid;  hydrochloric  acid,  HC1;  hydrobromic 
acid,  HBr;  hydriodic  acid,  HI;  hydrocyanic  acid,  HCX^;  hydrosul- 
phuric  acid,  H2S;  and  others. 

Group  B. — Acids  whose  radicals  are  precipitated  by  barium 
chloride,  from  a  neutral  solution:  sulphuric  acid,  H2SO4;  boric  acid, 
H3BO3;  sulphurous  acid,  H2SO3;  carbonic  acid,  H2CO3;  oxalic  acid, 


;  supurous  ac,  23;  caronc  a 
;  phosphoric  acid,  H^POi,  and  others, 
p  C. — Acids  whose  radicals  arc  not 

:  acetic  acid,  HC2H3O2;  nitric  acid,  II 


Group  C. — Acids  whose  radicals  arc  not  precipitated  in  groups  A 
and  B:  acetic  acid,  HC2H3O2;  nitric  acid,  IIXO;!;  and  others. 

GROUP  A. 

Acids  whose  radicals  are  precipitated  as  silver  salts  by  silver 
nitrate  in  the  presence  of  nitric  acid. 

The  principal  members  f  this  group  are  hydrochloric  aeid,  IIC1; 
hydrobromic  acid,  IIBr;  hydriodic  add,  III;  hydrocyanic  acid, 
IirX;  and  liydrosulphuric  acid.  II2S 

HYDROCHLORIC  ACID,  HC1. 

Muriatic  acid.  Spirit  of  salt. 

Hydrochloric  acid,  U.  S.  P.,  31-33%. 
Diluted  Hydrochloric  acid,  U.  S.  P.,  9.5-10.5%. 

Hydrochloric  acid  is  a  colorless  gas  with  a  sharp  suffocating  odor, 
which  is  exceedingly  soluble  in  water,  and  the  name  is  generally 
applied  to  the  solution.  The  acid  and  its  solution  are  irritant  or 
corrosive  poisons  when  inhaled  or  swallowed.  It  is  monobasic  and 
its  salts  are  called  chlorides,  formerly  muriates.  The  whole  acid 
combines  with  vegetable  alkaloids  and  such  salts  are  called  hydro- 
chlorides  or  hydrochlorates. 


HYDROCHLORIC  ACID  87 

The  chlorides  are  generally  colorless  and  most  of  them  are  soluble 
in  water.  Silver  chloride  and  mercurous  chloride  are  insoluble,  and 
lead  chloride  is  only  very  slightly  soluble;  also  the  basic  chlorides  of 
bismuth,  antimony  and  tin  are  insoluble  in  water. 

IMPORTANT  SALTS  OF  HYDROCHLORIC  ACID. 

Ammonium  chloride,  U.  S.  P.,  sal  ammoniac,  NH4C1. 

Barium  chloride,  BaCl2H2O. 

Calcium  chloride,  U.  S.  P.,  CaCl22H2O. 

Ferric  chloride,  U.  S.  P.,  FeCl36H2O. 

Gold  chloride,  AuCl3. 

Magnesium  chloride,  MgCl26H2O. 

Mercuric  chloride,  corrosive  mercuric  chloride,  U.  S.  P.,  corrosive 

sublimate,  HgCl2. 

Mercurous  chloride,  mild  mercurous  chloride,  U.  S.  P.,  calomel,  HgCl. 
Platinum  chloride,  PtCl4. 
Potassium  chloride,  KC1. 
Sodium  chloride,  U.  S.  P.,  salt,  NaCl. 
Stannic  chloride,  SnCl4. 
Stannous  chloride,  SnCl22H2O. 
Zinc  chloride,  U.  S.  P.,  ZnCl2. 

TESTS  FOR  SALTS  OF  HYDROCHLORIC  ACID. 

Use  separate  portions  of  a  solution  of  a;soluble  chloride,  as  XaCl, 
for  each  of  the  following  tests: 

1.  Add  AgNO3  T.S.  to  the  ^loride»blulion.     A  curdy  white 
precipitate  of  silver  chloride,  AgCl,  is  formed.     The  precipitate 
soon  darkens  upon  exposure  to  Jight  ancQp  insoluble  in  acids,  but  is 
readily  soluble  in  ammonjfl  wat  J*. 

2.  'Add  Pl)(('.Jl:{().j).j  T.S.  to  tlu»  chloride  solution.     A  white  pre- 
cipitate of  lead  chloride,  Pbrio,  is  formed  if  the  solution  is  not  too 
dilute.    The  precipitate  WslignHy  soluble  m  cold  water,  and  more 
soluble  in  hot  water. 

3.  Evaporate  a  portion  of  the  chloride  solution  to  dry  ness,  or 
nearly  so,  add  concentrated  H2SO4  and  heat.    HC1  gas  will  be  given 
off,  which  can  be  recognized  by  its  odor  and  by  the  white  fumes  of 
NH4C1  it  forms  if  a  glass  rod  or  stopper  wet  with  NH4OH  T.S.  is 
held  in  it. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.    NaCl  +  AgNOs  = 
HC1  +  AgN03  = 
CaCl2  +  AgNO3  = 


THE  ACIDS 


2.  KC1  +  Pb(C2H3O2)2  = 
HC1  +  Pb(N03)2  = 
BaCl2  +  Pb(C2H3O2)2  = 

3.  NaCl  +  H2SO4  = 
MgCl2  +  H2S04  - 
HC1  +  NH3  = 


HYDROBROMIC  ACID,  HBr. 

Diluted  hydrobromic  acid,  U.  S.  P.,  9.5-10.5%. 

Hydrobromic  acid  is  a  gas  which  resembles  hydrochloric  acid,  but 
is  less  stable  than  that  acid.  Its  salts  are  called  bromides. 

The  bromides  are  generally  colorless  and  soluble  in  water.  Silver 
bromide  and  mercurous  bromide  are  insoluble  and  lead  bromide  is 
only  slightly  soluble.  All  bromides  are  decomposed  by  chlorine 
and  bromine  is  set  free. 

IMPORTANT  SALTS  OF  HYDROBROMIC  ACID. 

Ammonium  bromide,  U.  S.  P.,  NH4Br. 
Calcium  bromide,  U.  S.  P.,  CaBr22H2O. 
Lithium  bromide,  U^B.  P.,  LiBr. 
Potassium  bromide,  YT.  S.  P.,  K 
Silver  bromide,  AgBr. 
Sodium  bromide,  U.  fi. 
Strontium  bromide,  if.  S«.,  Sr 


K 


TESTS  FOR  SALTS  OF  HYDROBROMIC  ACID. 

Use  a  separate  portion  of  a  solution  of  Bromide,  as  KBr,  for  each 
of  the  following  tests: 

1.  Add  AgNO3  T.Srto  the  br^ideAlution.     A  yellow-white 
precipitate  of  silver  bromide,  AgBr,  willbe  formed,  which  slowly 
darkens  upon  exposure  to  light.     The  precipitate  is  insoluble  in  nitric 
acid,  but  is  slowly  soluble  in  ammonia  water. 

2.  Add  Pb(C2H302)2  T.S.  to  the  bromide  solution.   $  A  white 
precipitate  of  lead  bromide,  PbBr2,  will  be  formed,  which  is  slightly 
soluble  in  water. 

3.  To  a  few  drops  of  the  bromide  solution  add  a  few  drops  of 
carbon  disulphide,  CS2>  and  slowly  add  chlorine  T.S.     Bromine 
will  be  set  free  and  will  dissolve  in  the  CS2  to  make  a  brown  solution. 
The  color  will  not  disappear  when  the  chlorine  is  in  excess. 


HYDRIODIC  ACID  89 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  KBr  +  AgNO8  = 
CaBr2  +  AgNO3  = 

2.  NaBr  +  Pb(C2H3O2)2  = 
NH4Br  +  Pb(N03)2  = 

3.  KBr  +  C12  = 
CaBr2  +  C12  = 

HYDRIODIC  ACID,  HI. 

Diluted  hydriodic  acid,  U.  S.  P.,  9.5-10.5  %. 

Hydriodic  acid  is  a  heavy  colorless  gas,  which  is  readily  soluble  in 
water.  It  is  unstable  and  readily  reacts  with  oxygen,  forming  water 
and  setting  iodine  free  when  oxidized.  The  salts  of  hydriodic  acid 
are  called  iodides. 

Silver  iodide,  mercurous  iodide,  mercuric  iodide  and  lead  iodide 
are  insoluble  in  water.  The  iodides  of  other  important  metals  are 
soluble.  All  iodides  are  decomposed  by  chlorine  or  bromine  and 
iodine  is  set  free. 

IMPORTANT  SALTS  OF  HYDRIODIC  ACID. 

Ammonium  iodide,  U.  S.  P.,  NH4I. 

Arsenous  iodide,  U.  S.  P.,  AsI3.   f 

Ferrous  iodide,  FeI2. 

Lead  iodide,  N.F.,  PbI2. 

Mercuric  iodide,  red  mercuric  iodide,  U.  jS.  P.,  biniodide  of  mercury, 

HgI2. 
Mercurous   iodide,  yellow  mercurous  iodide,  U.  S.  P.,  protoiodide 

of  mercury,  Hgl. 
Potassium  iodide,  U.  S.     .,  KI. 
Silver  iodide,  Agl. 
Sodium  iodide,  U.  S.  P., 
Strontium  iodide,  U.  S.  P.,  Sri. 
Zinc  iodide,  ZnI2. 

TESTS  FOR  SALTS  OF  HYDRIODIC  ACID. 

Use  a  separate  portion  of  a  solution  of  an  iodide,  as  KI,  for  each  of 
the  following  tests: 

1.  Add  AgNO3  T.S.  to  the  iodide  solution.    A  yellow  precipitate 
of  silver  iodide,  Agl,  will  be  obtained.    The  precipitate  slowly 
blackens  in  the  light  and  is  insoluble  in  HNO3  or  NH4OH. 

2.  Add  Pb(C2H3O2)2  T.S.  to  the  iodide  solution.    A  yellow  pre- 
cipitate of  lead  iodide,  PbI2,  will  be  obtained. 

3.  To  a  few  drops  of  the  iodide  solution  add  about  the  same 


90  THE  ACIDS 

amount  of  CS2,  and  slowly  add  chlorine  T.S.  Iodine  will  be  set  free 
and  will  dissolve  in  the  CS2  to  make  a  pink  solution.  When  the 
chlorine  is  in  excess  the  iodine  will  be  oxidized  to  iodic  acid,  HIO3, 
and  the  solution  will  be  decolorized. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  KI  +  AgN08  = 

NH4I  +AgC2H302  = 

2.  Nal  +  Pb(C2H302)2  = 
ZnI2  +  Pb(NO3)2  = 

3.  KI  +  C12  = 
CaI2  +  C12  = 

I2  +  C12  +  H20  =  HIO3  + 

HYDROCYANIC  ACID,  HCN. 

Prussic  acid. 
Diluted  hydrocyanic  acid,  U.  S,  P.,  1.9-2.1%. 

Hydrocyanic  acid  is  a  colorless  liquid  with  an  odor  resembling  that 
of  bitter  almonds.  It  is  soluble  in  water,  but  the  solution  slowly 
decomposes  through  the  influence  of  light  and  air.  The  salts  of 
hydrocyanic  acid  are  called  cyanides,  formerly  prussiates. 

The  cyanides  of  the  alkali  metals,  the  alkaline  earth  metals,  and 
mercuric  cyanide  are  soluble  in  water.  The  cyanides  of  other 
important  metals  are  insoluble.  The  acid  itself  and  all  of  its  salts 
are  deadly  poisons,  especially  dangerous  because  their  action  is  so 
rapid  that  there  is  little  time  for  the  administration  of  antidotes. 
Care  should  be  taken  not  to  inhale  the  gas  evolved  when  an  acid  is 
added  to  a  cyanide. 

• 
IMPORTANT  SALTS  OF  HYDROCYANIC  ACID. 

Mercuric  cyanide,  Hg(CN)2. 
Potassium  cyanide,  KCN. 
Silver  cyanide,  AgCN. 
Sodium  cyanide,  U.  S.  P.,  NaCN. 

TESTS  FOR  SALTS  OF  HYDROCYANIC  ACID. 

Use  a  separate  portion  of  a  solution  of  a  cyanide,  as  KCN,  for 
each  of  the  following  tests: 

1.  Slowly  add  AgNO3  T.S.  to  the  cyanide  solution.  A  white 
precipitate  of  silver  cyanide,  AgCN,  will  be  formed  when  the  AgNO3 
is  in  excess.  Add  an  excess  of  NH4OH  T.S.  The  precipitate  will 


HYDROSULPHURTC  ACID  91 

form  a  double  compound  and  dissolve.     Silver  cyanide  is  insoluble 
in  diluted  HNO3. 

2.  To  about  2  mils  of  the  cyanide  solution  in  a  test-tube,  add 
about  the  same  amount  of  each  of  FeSO4  T.S.  and  KOH  T.S.,  and 
warm,  when  potassium  ferrocyanide  K4Fe(CN)6  will  be  formed. 
Then  add  two  drops  of  FeCl3  T.S.  and  acidulate  with  HC1  T.S.    A 
blue  precipitate  of  ferric  ferrocyanide,  Fe4(Fe(CN)6)3,  will  be  formed. 

3.  To  about  2  mils  of  the  cyanide  solution  in  a  porcelain  evap- 
orating dish,  add  2  drops  of  yellow  ammonium  sulphide  and  heat 
on  a  water-bath  until  the  liquid  is  colorless,  potassium  sulphocyanate, 
KCNS,  being  formed.    Then  acidulate  with  HC1  T.S.  and  add  a 
drop  of  FeCl3  T.S.    A  red  coloration,  due  to  ferric  sulphocyanate, 
Fe(CNS)3,  will  be  produced. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  KCN  +  AgNO3  = 
HCN  +  AgN03  = 

2.  6KCN  +  FeSO4  =  K4Fe(CN)6  +  K2SO4 
FeCl3  +  K4Fe(CN)6  = 

3.  KCN  +  (NH4)2S2  =  KCNS  +  2NH3  +  H2S 
FeCl3  +  KCNS  = 

HYDROSULPHURIC  ACID,  H2S 

Hydrogen  sulphide.  Sulphuretted  hydrogen. 

Hydrosulphuric  acid  is  a  colorless  gas  having  a  characteristic  dis- 
agreeable odor;  it  is  moderately  soluble  in  water.  The  gas  in 
solution  readily  oxidizes  and  deposits  sulphur.  The  salts  of 
hydrosulphuric  acid  are  called  sulphides,  formerly  sulphurets. 
Many  metals  readily  combine  with  additional  atoms  of  sulphur  to 
form  polysulphides,  which  yield  hydrogen  sulphide  and  sulphur  on 
decomposition. 

The  sulphides  of  the  alkali  metals — potassium,  sodium,  lithium 
and  ammonium;  and  of  the  alkaline  earth  metals — magnesium, 
calcium,  strontium  and  barium;  are  soluble  in  water.  The  sulphides 
of  the  other  important  metals  are  insoluble. 

IMPORTANT  SALTS  OF  HYDROSULPHURIC  ACID. 

Ammonium  sulphide,  colorless  ammonium  sulphide,  (NH4)2S. 
Ammonium  poly  sulphide,  yellow  ammonium  sulphide,  (NH4)2SX. 
Antimonous  sulphide,  Sb2S3. 
Barium  sulphide,  BaS. 
Calcium  sulphide,  CaS. 


92  THE  ACIDS 

Ferrous  sulphide,  FeS. 

Lead  sulphide,  galena,  PbS. 

Mercuric  sulphide,  cinnabar,  vermilion,  HgS. 

Sodium  sulphide,  Na2S. 

Zinc  sulphide,  ZnS. 

TESTS  FOR  SALTS  OF  HYDRO  SULPHURIC  ACID. 

Use  a  separate  portion  of  a  solution  of  a  sulphide,  as  Na2S,  for  each 
of  the  following  tests: 

1.  Add  AgNO3  T.S.  to  the  sulphide  solution.    A  black  precipi- 
tate of  silver  sulphide,  Ag2S,  will  be  obtained,  which  is  insoluble  in 
diluted  acids. 

2.  Add  Pb(C2H3O2)2  T.S.   to  the  sulphide  solution.    A  black 
precipitate    of    lead    sulphide,    PbS,    will  be    obtained,    which    is 
insoluble  in  diluted  acids. 

3.  To  a  few  drops  of  the  sulphide  solution  add  H2SO4  T.S.,  and 
warm.    Hold  a  piece  of  filter  paper  moistened  with  Pb(C2H3O2)2 
T.S.  in  the  gas  evolved,  and  observe  the  odor.     Lead  sulphide,  PbS 
will  be  formed  and  will  blacken  the  paper,  and  the  odor  of  H2S  will 
be  observed. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  Na2S  +  AgN03  = 
CaS  +  AgNO3  = 

2.  Na2S  +  Pb('C2H3O2)2  = 
(NH4)2S  +  Pb(N03)2  = 

3.  K2S  +  H2SO4  = 
BaS  +  HC1  = 

H2S  +  Pb(C2H302)2  = 

SAMPLE  NUMBER  14. 

Analysis  of  a  solution  for  the  important  metals  and  for  group  A 

of  the  acids. 

The  sample  may  contain  chlorides,  bromides,  iodides,  cyanides  or 
sulphides  of  any  of  the  important  metals. 

A.  Examine  a  portion  of  the  sample  for  metals  by  the  tables  on 
pages  82-8.5. 

B.  To  about  5  mils  of  the  sample  add  H2SO4  T.S.  and  heat, 
observing  the  odor  of  the  vapor  given  off.     Hold  a  piece  of  lead 
acetate  test-paper  in  the*  vapor.     If  the  test-paper  is  not  blackened, 
sulphides  are  absent.     If  the  odor  of  H2S  is  perceived  and  the  test- 
paper  is  blackened,  sulphides  are  present. 


SULPHURIC  ACID  93 

C.  To  about  3  mils  of  the  sample  add  a  small  crystal  of  FeSO4  and 
an  excess  of  NaOH  T.S.  and  heat.  Then  add  two  drops  of  FeCl3 
T.S.  and  acidulate  with  HC1.  If  no  blue  precipitate  is  obtained, 
cyanides  are  absent.  A  blue  precipitate,  Fe4(Fe(CN)6)3,  shows  the 
presence  of  cyanides. 

Z).  To  a  few  drops  of  the  sample,  add  about  the  same  amount  of 
CS2  and  then  slowly  add  chlorine  T.S.  until  the  chlorine  is  in  excess. 
If  the  CS2  is  not  colored,  iodides  and  bromides  are  absent.  If  the 
€82  is  colored  pink  at  first  and  becomes  colorless  or  brown  when 
the  chlorine  is  in  excess,  iodides  are  present.  If  the  CS2  is  colored 
brown  when  the  chlorine  is  in  excess,  bromides  are  present. 

E.  If  sulphides,  cyanides,  iodides  and  bromides  are  absent, 
acidulate  a  small  portion  of  the  sample  with  HNO3  T.S.  and  add 
AgNO3  T.S.  If  no  precipitate  is  obtained,  chlorides  are  absent. 
A  white  precipitate  which  readily  dissolves  in  NH4OH  T.S.  shows 
the  presence  of  chlorides. 

If  sulphides  or  cyanides  are  present,  add  an  excess  of  HN03  T.S. 
to  a  small  portion  of  the  sample  and  boil,  under  a  hood,  to  expel 
H2S  or  HCN.  Then  add  AgNO3  T.S.  A  white  precipitate  shows 
the  presence  of  chlorides. 

If  bromides  or  iodides  are  present,  add  HNO3  T.S.  to  about  5  mils 
of  the  sample  and  boil  if  necessary  to  expel  H'2S  or  HCN.  Add  an 
excess  of  AgNO3  T.S.  and  filter.  Pass  about  2  mils  of  NH4OH  T.S. 
once  through  the  precipitate  on  the  filter.  Any  silver  chloride 
quickly  dissolves  as  ammonio-silver  chloride.  Acidulate  the  filtrate 
with  HNO3.  A  white  precipitate  shows  the  presence  of  chlorides. 
This  test  effectually  separates  chlorides  from  iodides,  but  the 
separation  from  bromides  is  not  complete,  as  silver  bromide  slowly 
dissolves  in  ammonia  water.  To  completely  separate  chlorides  from 
bromides  is  difficult  of  accomplishment,  and  the  methods  used  are 
beyond  the  scope  of  this  manual. 

GROUP  B. 

Acids  whose  radicals  are  precipitated  from  a  neutral  solution  as 
barium  salts  by  barium  chloride  or  barium  nitrate. 

The  principal  members  of  this  group  are  sulphuric  acid,  H2SO4; 
sulphurous  acid,  H2SO3;  carbonic  acid,  H2CO3;  oxalic  acid,  H2C2O4; 
and  phosphoric  acid,  H3PO4. 

SULPHURIC  AGD,  H2SO4. 

Oil  of  vitrol. 

Sulphuric  acid,  U.S.P.,  93-95%.  Diluted  sulphuric  acid,  U.S.P.  9.5-10.5%. 
Sulphuric  acid  is  a  heavy,  colorless  and  odorless  liquid  which  is 
hygroscopic  and  non-volatile  at  ordinary  temperatures,  but  may  be 


94  THE  ACIDS 

distilled  unchanged.  The  concentrated  acid  attacks  and  destroys 
organic  matter  and  inflicts  serious  burns  if  it  conies  in  contact  with 
the  skin.  It  has  a  strong  affinity  for  water  and  much  heat  is  gen- 
erated when  the  strong  acid  and  water  are  mixed.  On  this  account, 
the  acid  should  always  be  poured  into  the  water,  when  they  are 
being  mixed,  so  that  the  rise  in  temperature  will  be  gradual. 

Sulphuric  acid,  being  dibasic,  forms  normal  salts  called  sulphates, 
and  acid  salts  called  bisulphates.  The  sulphates  and  bisulphates 
are  generally  colorless  and  soluble  in  water.  Lead  sulphate,  barium 
sulphate  and  strontium  sulphate  are  insoluble,  and  mercurous 
sulphate  and  calcium  sulphate  are  but  slightly  soluble. 

IMPORTANT  SALTS  OF  SULPHURIC  ACID. 

Aluminum  and  ammonium  sulphate,  ammonia  alum,  A1NH4(SO4)2- 

12H20. 

Aluminum  and  potassium  sulphate,  potash  alum,  A1K(SO4)212H20. 
Aluminum  sulphate,  A12(SO4)318H2O. 
Ammonium  sulphate,  (NH4)?SO4. 

Barium  sulphate,  heavy  spar,  permanent  white,  BaSO4. 
Cadmium  sulphate,  CdSO4. 
Calcium  sulphate,  gypsum,  CaSO42H2O. 
Calcium  sulphate,  anhydrous,  plaster  of  Paris,  CaS04. 
Chromium  potassium  sulphate,  chrome  alum,  CrK(S04)212H2O. 
Cobalt  sulphate,  CoSO47H2O. 

Copper  sulphate,  blue  vitriol,  blue  stone,  CuSO45H2O. 
Ferric  ammonium  sulphate,  ferric  alum,  FeNH4(SO4)212H2O. 
Ferric  subsulphate,  Monsel's  salt,  Fe4O(S04)5. 
Ferric  sulphate,  iron  tersulphate,  Fe2(SO4)3. 
Ferrous  sulphate,  green  vitriol,  copperas,  FeSO47H20. 
Magnesium  sulphate,  Epsom  salt,  MgS047H2O. 
Manganous  sulphate,  MnSO44H2O. 
Mercuric  sulphate,  HgSO4. 
Mercurous  sulphate,  Hg2SO4. 
Nickel  sulphate,  NiSO4. 
Potassium  sulphate,  K2SO4. 
Silver  sulphate,  Ag2SO4. 

Sodium  sulphate,  Glauber's  salt,  Na2SO410H2O. 
Strontium  sulphate,  SrSO4. 
Zinc  sulphate,  white  vitriol,  ZnSO47H2O. 

TESTS  FOR  SALTS  OF  SULPHURIC  ACID. 

Use  a  separate  portion  of  a  solution  of  a  sulphate,  as  Na2SO4,  for 
each  of  the  following  tests: 

1.    Add  BaCl2  T.S.  to  the  sulphate  solution.    A  finely  divided 


SULPHUROUS  ACID  95 

white  precipitate  of  barium  sulphate,  BaSO4,  will  be  obtained,  which 
is  insoluble  in  HC1  or  HNOs. 

2.  Add  Pb(C2H302)2  T.S.  to  the  sulphate  solution.  A  white 
precipitate  of  lead  sulphate,  PbSO4,  will  be  obtained,  which  is 
insoluble  in  acids. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  NasSCX  +  BaCl2  = 
MgS04  +  Ba(N03)2  = 

2.  K2SO4  +  Pb(C2H3O2)2  = 
ZnSO4  +  Pb(NO3)2  = 

SULPHUROUS  ACID,  H2SO3. 

Sulphurous  acid  is  formed  when  its  anhydride,  SO2,  is  dissolved  in 
water,  but  the  acid  is  so  unstable  that  it  cannot  be  isolated. 

The  acid,  being  dibasic,  forms  normal  salts,  called  sulphites,  and 
acid  salts,  called  bisulphites.  The  sulphites  and  bisulphites  of  the 
alkali  metals  and  the  bisulphites  of  the  alkaline  earth  metals  are 
soluble  in  water.  The  other  metals  do  not  generally  form  bisulphites 
and  their  sulphites  are  insoluble.  Sulphites  and  bisulphites  are 
decomposed  by  most  acids,  the  acid  anhydride,  sulphur  dioxide, 
SO2,  being  generated. 

IMPORTANT  SALTS  .OF  SULPHUROUS  ACID. 

Calcium  sulphite,  CaSO32H2O. 

Calcium  bisulphite,  bisulphite  of  lime,  Ca(HS03)2. 

Sodium  sulphite,  Na2SO37H2O. 

Exsiccated  sodium  sulphite,  U.  S.  P.,  Na2SO3. 

Sodium  bisulphite,  NaHSO3. 

TESTS  FOR  SALTS  OF  SULPHUROUS  ACID. 

Use  a  separate  portion  of  a  solution  of  a  sulphite,  as  Na2S03,  for 
each  of  the  following  tests: 

1.  Add  BaCl2  T.S.  to  the  sulphite  solution,  followed  by  an  excess 
of  HC1  and  then  by  a  few  drops  of  nitrohydrochloric  acid.    A  white 
precipitate  of  barium  sulphite,  BaSO3,  will  be  obtained,  which  will 
dissolve  in  the-  HC1;  and  a  precipitate  of  barium  sulphate,  BaSO4, 
will  be  formed  upon  adding  the  nitrohydrochloric  acid. 

2.  Add  H2SO4  T.S.  to  a  small  quantity  of  the  sulphite  solution 
and  heat  the  liquid,  observing  the  odor  of  the  gas  evolved.    The 
characteristic  choking  odor  is  that  of  sulphur  dioxide,  SO2.    The 
bleaching  effect  of  the  gas  on  organic  coloring  matter  may  also  be 
obtained,  most  easily  by  holding  a  colored  flower,  as  a  violet,  in  the 
gas. 


96  THE  ACIDS 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 


1.  NasSOs  +  BaCl2  = 
Ca(HS03)2  +  Ba(NO3)2  = 

2.  Na^SOs  +  H2SO4  = 
NaHS03  +  HC1  = 

CARBONIC  ACID  H2CO3. 

Carbonic  acid  appears  to  be  formed  when  carbon  dioxide  is  dis- 
solved in  water,  as  in  "soda  water."  The  compound,  however,  is 
so  unstable  that  it  cannot  be  isolated.  The  acid  forms  normal 
salts  called  carbonates,  and  acid  salts  called  bicarbonates.  Carbon- 
ates and  bicarbonates  are  decomposed  by  nearly  all  other  acids, 
yielding  water  and  carbon  dioxide.  The  carbonates  of  the  alkali 
metals  are  soluble  in  water.  All  other  carbonates  are  insoluble  in 
water,  though  several  of  them  will  form  unstable  bicarbonates  and 
dissolve  in  water  containing  carbon  dioxide  in  solution. 

IMPORTANT  SALTS  OF  CARBONIC  ACID. 

Ammonium  carbonate,  U.  S.  P.,  ammonium  bicarbonate  +  ammo- 

nium carbamate,  NH4HCO3  +  NH4NH2CO2. 
Calcium    carbonate,    limestone,    marble,    chalk,  prepared   chalk, 

U.  S.  P.,  precipitated  calcium  carbonate,  U.  S.  P.,  CaCO3. 
Ferrous  carbonate,  FeCO3. 

Lead  sub-carbonate,  white  lead,  (PbCO3)2Pb(OH)2. 
Lithium  carbonate,  U.  S.  P.,  Li2CO3. 
Magnesium  carbonate,  magnesite,  MgCO3. 
Magnesium  carbonate,  U.  S.  P.,  magnesia  alba,  (MgCO3)4Mg(OH)2- 

5H2O. 

Potassium  carbonate,  U.  S.  P.,  potash,  salt  of  tartar,  K2CO3. 
Potassium  bicarbonate,  U.  S.  P.,  saleratus,  KHCO3. 
Sodium  carbonate,  soda,  washing  soda,  Na2CO310H20. 
Monohydrated  sodium  carbonate,  U.  S.  P.,  Na2CO3H2O. 
Sodium  bicarbonate,  U.  S.  P.,  baking  soda,  NaHCO3. 
Zinc  carbonate,  calamine,  ZnCO3. 

TESTS  FOR  SALTS  OF  CARBONIC  ACID. 

Use  a  separate  portion  of  a  solution  of  a  carbonate,  as  Na2CO3,  for 
each  of  the  following  tests  : 

1.  Add  BaCl2  T.S.  to  the  carbonate  solution.    A  white  precipi- 
tate of  barium  carbonate,  BaCO3,  will  be  formed,  which  is  soluble  in 
acids. 

2.  Add  -CaCl2  T.S.  to  the  carbonate  solution.    A  white  precipi- 
tate of  calcium  carbonate,  CaCO3,  will  be  formed,  which  is  soluble 
in  acids. 


OXALIC  ACID  97 

3.  To  a  portion  of  the  carbonate  solution  contained  in  a  test- 
tube  provided  with  a  delivery  tube,  add  H2SO4  T.S.  and  conduct  the 
gas  generated  into  lime  water,  Ca(OH)2.  The  carbonate  will  be 
decomposed,  yielding  carbon  dioxide,  CO2,  and  this  will  precipitate 
calcium  carbonate,  CaCO3,  from  the  lime  water. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  Na2CO3  +  BaCl2  = 
KHC03  +  Ba(N03)2  = 

2.  K2CO3  +  CaCl2  = 
(NH4)2C03  +  CaS04  = 

3.  NaaCOs  +  H2SO4  = 
NaHCOa  +  HCL  = 
Ca(OH)2  +  C02  = 

SAMPLE  NUMBER  15. 

Analysis  of  a  solution  for  the  important  metals  and  for  the  radicals 
of  sulphuric,  sulphurous  and  carbonic  acids. 

The  sample  may  contain  any  of  the  important  metals,  as  sul- 
phates, sulphites  or  carbonates. 

A.  Examine  a  portion  of  the  sample  for  metals  by  the  tables  on 
pages  82-85. 

B.  To  a  portion  of  the  sample,  add  HC1  until  the  reaction  is  acid 
and  then  add  BaCl2  T.S.     If  no  precipitate  is  produced  sulphates 
are  absent.     A  white  precipitate  shows  the  presence  of  sulphates. 

C.  To  a  portion  of  the  sample,  contained  in  a  test-tube  provided 
with  a  delivery  tube,  add  H2SO4  T.S.,  and  heat.     If  a  gas  is  evolved, 
observe  its  odor  and  pass  it  through  lime  water.     If  there  is  no 
effervescence  of  an  odorous  gas  and  the  lime  water  is  not  rendered 
turbid,  sulphites  and  carbonates  are  absent.     If  there  is  an  effer- 
vescence of  a  gas  having  the  odor  of  burning  sulphur,  sulphites  are 
present.     If  there  is  an  effervescence  of  an  odorless  gas  which  renders 
lime  water  turbid,  carbonates  are  present.     If  the  odor  of  burning 
sulphur  is  observed  and  the  lime  water  is  rendered  turbid,  both 

sulphites  and  carbonates  are  present. 

i 

OXALIC  ACID,  H2C2O42H2O. 

Oxalic  acid  is  an  organic  acid,  whose  salts  are  frequently  used  as 
reagents.  It  is  a  soluble  crystalline  solid.  The  acid  is  dibasic  and 
forms  normal  and  acid  salts,  called  oxalates,  which  are  generally 
colorless.  The  acid  and  its  salts  are  poisonous.  The  oxalates  of 
the  alkali  metals  and  of  chromium  are  soluble  in  water.  The 
7 


98  THE  ACIDS 

oxalates  of  all  other  important  metals  are  insoluble  in  water,  but 
will  dissolve  in  mineral  acids,  if  the  metals  present  form  soluble 
salts  with  the  acids. 

IMPORTANT  SALTS  OF  OXALIC  ACID. 


Ammonium  oxalate, 

Calcium  oxalate,  CaC2O4H2O. 

Cerium  oxalate,'  Ce2(C2O4)310H2O. 

Potassium  acid  oxalate,  salt  of  lemon,  KHC2O42H2O. 

Sodium  oxalate,  Na2C2C>4. 

TESTS  FOR  SALTS  OF  OXALIC  ACID. 

Use  a  separate  portion  of  a  solution  of  an  oxalate,  as  (NH4)2C204 
for  each  of  the  following  tests  : 

1.  Add  BaCl2  T.S.  to  the  oxalate  solution.  A  white  precipitate  oj 
barium  oxalate,  BaC2O4,  will  be  formed;  the  precipitate  is  soluble  ir 
acids  which  form  soluble  salts  with  barium,  including  acetic  acid. 

2.  Add  CaCl2  T.S.  to  the  oxalate  solution.    A  white  precipitate 
of  calcium  oxalate,  CaC2O4,  will  be  formed,  which  is  soluble  ir 
mineral  acids  but  not  in  acetic  acid. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  (NH4)2C2O4  +  BaCl2  = 
Na2C2O4  +  Ba(N03)2  = 

2.  K2C204  +  CaCl2  = 
KHC2O4  +  CaSO4  = 

PHOSPHORIC  ACID,  H3PO4. 

Phosphoric  acid,  U.S.P.,  85-88%.      Dilute  phosphoric  acid,  U.S.P.,  9.5-10.5% 
Orthophosphoric  acid. 

Orthophosphoric  acid  is  a  colorless  deliquescent  crystalline  solic 
commonly  seen  as  a  syrupy  solution  in  water  and  whose  salts  are  the 
phosphates.  It  is  not  volatile  but  when  heated  loses  water  anc 
changes,  first  to  pyrophosphoric  acid,  H4P2O7,  and  then  to  meta- 
phosphoric  acid,  HPOs. 

Orthophosphoric  acid  is  tri-basic  and  forms  normal  salts  and  twc 
series  of  acid  salts.  Salts  in  which  only  two  of  the  three  hydroger 
atoms  are  replaced  by  metals,  as  Na2HP04,  are  common,  and  are 
frequently  called  phosphates.  The  phosphates  of  the  alkali  metah 
and  the  acid  phosphates  of  a  few  other  metals  are  soluble  in  water 
All  other  phosphates  are  insoluble. 


PHOSPHORIC  ACID  99 

IMPORTANT  SALTS  OF  PHOSPHORIC  ACID. 

Ammonium  phosphate,  (NH4)2HP04. 

Calcium  phosphate,  Ca3(PO4)2. 

Ferric  phosphate,  FePO4. 

Sodium  phosphate,  U.  S.  P.,  Na2HP0412H2O. 

TESTS  FOR  SALTS  OF  PHOSPHORIC  ACID. 

Use  a  separate  portion  of  a  solution  of  a  phosphate,  as  Na2HPO4, 
for  each  of  the  following  tests: 

1.  Add  BaCl2  T.S.  to  the  phosphate  solution.    A  white  precipi- 
tate of  barium  phosphate,  BaHPO4,  will  form,  which  is  soluble  in 
acids. 

2.  To  the  phosphate  solution  add  CaCl2  T.S.,  followed  by  an 
excess  of  HC2H3O2  T.S.    A  white  precipitate  of  calcium  phosphate, 
CaHPO4,  will  be  obtained,  which  will  dissolve  in  the  acetic  acid. 

3.  To  the  phosphate  solution  add  NH4C1  T.S.,  NH4OH  T.S.  and 
MgS04  T.S.  in  the  order  given.    A  white  precipitate  of  magnesium 
ammonium  phosphate,  MgNH4PO4,  will  form,  which  is  soluble  in 
acids. 

4.  To  about  2  mils  of  (NH4)2MoO4  T.S.  add  a  few  drops  of  the 
phosphate   solution   and  let   stand.    A  yellow  precipitate  of  am- 
monium phosphomolybdate,  (NH4)3PO4(Mo03)]2,  will  form  slowly. 
The  precipitate  is  soluble  in  alkalis. 

COMPLETE  AND  BALANCE  *HE  FOLLOWING  EQUATIONS: 

1.  Na2HP04  +  BaCl2  = 
(NH4)2HP04  +  Ba(N03)2  = 

2.  K2HP04  +  CaCl2  = 

CaHPO4  +  HC2H3O2  =  Ca(H2P04)2  + 

3.  NaHP04  +  NH4OH  +  MgSO4  = 
(NH4)2HPO4  +  MgCl2  = 

4.  Na,HPO4  +  12(NH4)2MoO4  +  23HNO3 

=  (NH4)3P04(Mo03)12  +  21NH4NO3  +  2NaNO3  +  12H2O. 

SAMPLE  NUMBER  16. 

Analysis  of  a  solution  for  the  principal  metals  and  for  the  radicals 
of  oxalic  and  phosphoric  acids. 

The  sample  may  contain  any  of  the  important  metals  as  oxalates 
or  phosphates. 

A.  Add  NaOH  T.S.  to  a  portion  of  the  sample  until  the  reaction 
is  alkaline,  then  HC2H302  T.S.  until  the  reaction  is  acid,  then  CaCl2 
T.S.  and  let  stand.  If  no  precipitate  is  produced  oxalic  acid  is 


100  THE  ACIDS 

absent.    A  white  precipitate,  which  may  be  formed  slowly,  shows 
the  presence  of  oxalates. 

B.  To  about  3  mils  of  (NH4)2MoO4  T.S.  in  a  test-tube  add  a  few 
drops  of  the  sample  and  let  stand.     If  no  precipitate  is  formed 
phosphates  are  absent.     A  yellow  precipitate,  slowly  formed,  shows 
the  presence  of  phosphates. 

About  10  mils  of  the  sample  should  be  examined  for  the  metals 
by  the  tables  on  pages  82-85,  with  the  following  changes  if  oxalates 
or  phosphates,  or  both,  are  present. 

C.  If  oxalates  are  present  evaporate  the  filtrate  from  the  second 
group  to  a  small  volume,  transfer  to  a  porcelain  crucible,  add 
concentrated  HNO3  and  gently  ignite,  which  treatment  will  destroy 
oxalic  acid.     Then  dissolve  the  residue  in  HC1  T.S.  and  proceed  with 
the  examination  for  group  3  of  the  metals. 

D.  If  phosphates  are  present  concentrate  the  filtrate  from  the 
second  group  of  metals  to  about  5  mils,  or  use  the  solution  obtained 
in  C  after  destroying  oxalic  acid,  if  present.     Mix  with  three  times 
its  volume  of  (NH4)2MoO4  T.S.,  warm  the  mixture  and  let  it  stand 
for  some  time.     Filter  and  test  the  filtrate.     Evaporate  the  filtrate 
nearly  to  dryness  to  expel  nitric  acid.     Dilute  with  about  10  mils 
of  water,  add  a  few  drops  of  HC1  and  pass  H2S  through  the  solution 
to  precipitate  the  excess  of  molybdenum,  then  filter  and  proceed 
with  the  examination  for  group  3  of  the  metals,  etc. 

GROUP  C. 

Acids  whose  radicals  are  not  precipitated  by  any  group  reagent. 
The  principal  members  of  this  group  are  boric  acid,  H3BO3;  acetic 
acid,  HC2H3O2;  and  nitric  acid,  HNO3. 

BORIC  ACID,  H3B03. 

Boric  acid,  U.  S.  P. 
Boracic  acid.  Orthoboric  acid. 

Boric  acid  is  a  crystalline  solid  which  is  moderately  soluble  in 
water.  The  acid  decomposes  when  heated,  forming  successively 
metaboric  acid,  HBO2;  pyroboric  acid  or  tetraboric  acid,  H2B4O?; 
and  finally  the  acid  anhydride,  boron  trioxide,  B2O3.  The  salts  of 
ortho-boric  acid  are  unstable  and  seldom  seen.  The  borates, 
generally,  are  the  salts  of  tetraboric  acid  or  metaboric  acid,  but 
when  decomposed  in  the  presence  of  water  they  yield  orthoboric 
acid.  The  borates  are  decomposed  by  nearly  all  other  acids.  The 
borates  of  the  alkali  metals  are  soluble  in  water.  The  borates  of 
other  metals  are  insoluble  in  water  but  many  of  them  become 
soluble  in  the  presence  of  an  excess  of  boric  acid. 


ACETIC  ACID  101 

IMPORTANT  SALTS  OF  BORIC  ACID. 

Glyceryl  borate,  boroglycerin,  C3H5BO3. 

Sodium  borate,  U.  S.  P.,  sodium  tetraborate,  sodium  pyroborate, 
borax, 


TESTS  FOR  SALTS  OF  BORIC  ACID. 

Use  a  separate  portion  of  a  solution  of  a  borate,  as  Na^O?,  for 
each  of  the  following  tests: 

1.  Add  BaCl2  T.S.  to  the  borate  solution.    A  white  precipitate  of 
barium  metaborate,  Ba(BO2)2,  will  form  if  the  solution  is  not  too 
dilute.    The  precipitate  is  soluble  in  an  excess  of  the  reagent  and 
in  HNO3  or  HCL 

2.  Add  H2SO4  T.S.  to  a  portion  of  borate  solution  in  a  porcelain 
evaporating  dish;  then  add  about  twice  the  volume  of  alcohol,  stir 
and  ignite.    The  outer  mantle  of  the  flame  will  show  a  green  color. 
Before  applying  this  test  it  is  necessary  to  remove  copper  and  barium 
from  the  solution,  if  they  are  present,  by  H2S  and  H2SO4,  respectively. 

3.  Add  HC1  T.S.  to  the  borate  solution,  dip  a  piece  of  turmeric 
paper  in  the  liquid  and  dry  the  paper.    The  paper  will  be  colored 
red-brown,  changing  to  red  upon  drying. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  Na2B407  +  BaCl2  +  H2O  =  Ba(B02)2  + 
Na2B4O7  +  Ba(NO3)2  +  H^O  = 

2.  Na2B4O7  +  H2SO4  +  H2O  =  H3BO3  + 

3.  Na2B4O7  +  HC1  +  H2O  =  H3B03  + 

ACETIC  ACID,  HC2H3O2. 

Acetic  acid,  U.  S.  P.,  36-37%. 
Diluted  acetic  acid,  U.S.P.,  5.7-6.3%.     Glacial  acetic  acid,  U.S.P.,  99+%. 

Acetic  acid  is  a  colorless  liquid  with  a  characteristic  pungent  odor. 
The  pure  acid  corrodes  the  skin  and  freezes  to  a  glass-like  solid  at 
temperatures  below  15°.  The  diluted  acid  has  about  the  strength 
of  vinegar,  from  which  it  was  first  prepared. 

The  acid  is  monobasic  and  its  salts  are  called  acetates.  All  the 
normal  acetates  are  soluble  in  water,  but  insoluble  basic  acetates  are 
formed  by  several  of  the  metals. 

IMPORTANT  SALTS  OF  ACETIC  ACID. 

Ammonium  acetate,  NH4C2H3O2. 
Copper  acetate,  Cu(C2H3O2)2H2O. 
Copper  aceto-arsenite,  paris  green,  Cu(C2H3O2)2  +  3Cu(As02)2. 


102  THE  ACIDS 

Copper  subacetate,  verdigris,  Cu20(C2H302)2. 

Ferric  acetate,  Fe(C2H3O2)3. 

Lead  acetate,  U.  S.  P.,  sugar  of  lead,  Pb(C2H3O2)23H2O. 

Lead  sub-acetate,  Pb2O(C2H3O2)2. 

Potassium  acetate,  U.  S.  P.,  KC2H3O2. 

Sodium  acetate,  U.  S.  P.,  NaC2H3O23H2O. 

Zinc  acetate,  U.  S.  P.,  Zn(C2H3O2)22H2O. 

TESTS  FOR  SALTS  OF  ACETIC  ACID. 

Use  a  separate  portion  of  a  solution  of  an  acetate,  as  NaC2H302, 
for  each  of  the  following  tests: 

1.  To  a  portion  of  the  acetate  solution  add  H2S04  T.S.,  boil  the 
mixture  and  observe  the  odor  of  the  vapor,  which  will  be  that  of  the 
acetic  acid,  HC2H3O2,  set  free. 

2.  To  about  2  mils  of  the  acetate  solution,  add  a  few  drops  of 
alcohol,  C2H5OH,  and  about  1  mil  of  H2SO4  T.S.  and  heat,  observing 
the  odor  of  the  vapor.     The  characteristic  fruity  odor  of  ethyl 
acetate,  C2H5C2H3O2  will  be  observed. 

3.  Add  FeCl3  T.S.  to  the  acetate  solution  which  should  be  neutral 
or  only  slightly  acid  in  reaction,  and  boil  the  mixture.    Ferric  ace- 
tate, Fe(C2H3O2)3,  is  formed  and  colors  the  liquid  red.     Upon  boiling 
a  brown-red  precipitate  of  a  basic  ferric  acetate,  FeOC2H3O2,  will  be 
obtained. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  NaC2H302  +  H2S04  = 
Pb(C2H3O2)2  +  HC1  = 

2.  NaC2H3O2  +  H2S04  +  C2H5OH  =  NaHSO4  + 

3.  KC2H302  +  FeCl3  = 
Zn(C2H302)2  +  FeCl3  = 
Fe(C2H302)3  +  H20  = 

NITRIC  ACID,  HNO3. 

Aqua  fortis. 
Nitric  acid,  U.S.P.,  67-69%. 

Nitric  acid  is  a  colorless  fuming  liquid  which  is  unstable  and 
difficult  to  prepare  in  a  pure  state.  The  official  acid  is  a  concen- 
trated solution  and  is  a  dangerously  corrosive  and  poisonous  liquid. 
It  is  a  strong  acid  and  also  a  powerful  oxidizing  agent.  "Fuming 
nitric  acid"  is  brown  in  color  and  is  a  solution  of  nitrogen  peroxide, 
NO2,  in  concentrated  nitric  acid. 

The  acid  is  monobasic  and  its  salts  are  called  nitrates.  All 
normal  nitrates  are  soluble  in  water. 


NITRIC  ACID  103 

IMPORTANT  SALTS  OF  NITRIC  ACID. 

Ammonium  nitrate,  NH4NO3. 

Barium  nitrate,  Ba(NO3)2. 

Bismuth  nitrate,  Bi(NO3)3. 

Bismuth  sub-nitrate,  U.  S.  P.,  approximately  Bi(OH)2NO3. 

Ferric  nitrate,  Fe(NO3)3. 

Lead  nitrate,  Pb(NO3)2. 

Mercuric  nitrate,  Hg(N03)24H2O. 

Potassium  nitrate,  U.  S.  P.,  saltpetre,  nitre,  KN03. 

Silver  nitrate,  U.  S.  P.,  lunar  caustic,  AgNO3. 

Sodium  nitrate,  Chili  saltpetre,  NaNO3. 

Strontium  nitrate,  Sr(NO3)2. 

TESTS  FOR  SALTS  OF  NITRIC  ACID. 

Use  a  separate  portion  of  a  solution  of  a  nitrate,  as  KNO3,  for  each 
of  the  following  tests: 

1.  To  a  portion  of  the  nitrate  solution  add  an  equal  volume  of 
H2S04  T.S.  and  a  small  amount  of  metallic  copper  and  heat  the 
mixture.  The  nitrate  is  decomposed  and  gaseous  nitric  oxide,  NO, 
is  formed.  This  oxidizes  in  the  air  and  the  brown  gas,  nitrogen 
peroxide,  N02,  is  obtained. 

2  BROWN  RING  TEST.  Mix  about  1  mil  of  the  nitrate  solution 
with  about  the  same  volume  of  concentrated  H2SO4,  and  cool  the 
mixture.  Then  carefully  pour  some  freshly  made  FeSO4  T.S.  down 
the  side  of  the  test-tube  so  that  the  liquids  do  not  mix  and  let  stand. 
The  liberated  nitric  acid  will  oxidize  some  of  the  iron  to  form  ferric 
sulphate  and  this  will  give  a  brown  ring  between  the  two  layers. 

The  test  may  also  be  applied  by  mixing  the  solution  to  be  tested 
with  FeSO4  T.S.  and  superimposing  the  mixture  on  concentrated 
H2SO4  in  a  test-tube. 

COMPLETE  AND  BALANCE  THE  FOLLOWING  EQUATIONS: 

1.  2KN03  +  4H2SO4  +  3Cu  =  K2SO4  +  3CuS04  +  4H2O  +  2NO 
2NO  +  O2  =  2NO2 

2.  KNO3  +  H2SO4  = 

6FeS04  +  3H2SO4  +  2HNO3  =  3Fe2(SO4)3  +  4H20  +  2NO 

SAMPLE  NUMBER  17. 

Analysis  of  a  solution  for  the  principal  metals  and  for  the  radicals 
of  boric,  acetic  and  nitric  acids. 

The  sample  may  contain  any  of  the  important  metals,  as  borates, 
acetates  or  nitrates. 


104  THE  ACIDS 

A.  Examine  a  portion  of  the  sample  for  metals  by  the  tables  on 
pages  82-85. 

B.  To  a  portion  of  the  sample  add  HC1  T.S.,  dip  a  piece  of 
turmeric  paper  in  the  liquid  and  dry  the  paper.     If  the  color  of  the 
paper  is  unchanged  borates  are  absent.     If  the  paper  is  turned  red- 
brown,  changing  to  dark  red  upon  drying,  borates  are  present. 

C.  To  a  portion  of  the  sample  add  a  few  drops  of  alcohol  followed 
by  an  excess  of  H2SO4  T.S.,  heat  the  liquid  and  observe  the  odor  of 
its  vapor.     If  no  odor  of.  ethyl  acetate  can  be  observed,  acetates  are 
absent.     If  the  odor  of  ethyl  acetate  is  observed,  acetates  are  present. 

D.  Mix  some  of  the  sample  with  about  the  same  volume  of  con- 
centrated H2SO4  in  a  test-tube,  cool  the  liquid,  superimpose  FeS04 
T.S.  on  it  and  let  stand.     If  no  brown  ring  forms  at  the  junction  of 
the  two  liquids,  nitrates  are  absent.     If  a  brown  ring  forms  at  the 
junction  of  the  liquids,  nitrates  are  present. 

SAMPLE  NUMBER  18. 
Analysis  of  a  solution  for  the  important  metals  and  acids. 

The  sample  may  contain  any  of  the  important  metals  as  salts 
of  any  of  the  acids  studied. 

A.  Add  NaOH  T.S.  to  a  portion  of  the  sample  until  the  reaction 
is  alkaline,  then  HC2H3O2  T.S.  until  the  reaction  is  acid,  then  CaCl2 
T.S.  and  let  stand.     If  no   precipitate  is  produced,  oxalates  are 
absent.     A  white  precipitate,  which  may  be  formed  slowly,  shows 
the  presence  of  oxalates. 

B.  To  about  3  mils  of  (NH4)2MoO4  T.S.  in  a  test-tube  add  a  few 
drops  of  the  sample,  warm  and  let  stand.     If  no  precipitate  forms, 
phosphates  are  absent.     A  yellow  precipitate,  slowly  formed,  shows 
the  presence  of  phosphates. 

C.  Examine  a  portion  of  the  sample  for  metals  by  the  tables  on 
pages  82-85,  removing  oxalates  and  phosphates  if  present  as  directed 
on  pages  99-100. 

D.  To  about  5  mils  of  the  sample  add  an  excess  of  H2SO4  T.S. 
and  heat.     If  a  gas  is  given  off  observe  its  odor,  hold  a  piece  of  paper 
wet  with  Pb(C2H3O2)2  T.S.  in  the  gas  and  pass  some  of  it  through 
lime  water. 


A  gas  having  the  odor  i     A  gas  having  the  odor  of 
of  burning  sulphur  shows    HsS  and  which  blackens 


the  presence  of  sulphites.    Pb^zRsO^z     shows    the 
presence  of  sulphides. 


An  odorless  gas  which 
renders  lime  water  tur- 
bid shows  the  presence 
of  carbonates. 


E.    Add  a  slight  excess  of  HNO3  T.S.  to  a  portion  of  the  sample, 
followed  by  AgNO3  T.S.     If  no  precipitate  is  produced,  the  acids  of 


NITRIC  ACID  105 

group  A  are  absent:  pass  to  paragraph  I.  If  a  precipitate  is  pro- 
duced, proceed  by  paragraph  F,  remembering  that  sulphides,  if 
present,  will  have  been  detected  in  D. 

F.  To  about  3  mils  of  the  sample  add  a  small  crystal  of  FeSO4  and 
an  excess  of  NaOH  T.S.  and  heat.    Then  add  two  drops  of  FeCl3 
T.S.  and  acidulate  with  HC1.    A  blue  precipitate,  Fe4(Fe(CH)6)3, 
shows  the  presence  of  cyanides. 

G.  To  a  few  drops  of  the  sample,  add  about  the  same  amount  of 
€82  and  then  slowly  add  chlorine  T.S.  until  the  chlorine  is  in  excess. 
If  the  CS2  is  not  colored,  iodides  and  bromides  are  absent.     If  the 
CS2  is  colored  pink  at  first  and  becomes  colorless  or  brown  when  the 
chlorine  is  in  excess,  iodides  are  present.     If  the  CS2  is  colored 
brown  when  the  chlorine  is  in  excess,  bromides  are  present. 

H.  If  sulphides,  cyanides,  iodides  and  bromides  are  absent, 
acidulate  a  small  portion  of  the  sample  with  HNO3  T.S.  and  add 
AgX03  T.S.  A  white  precipitate  which  readily  dissolves  in  NH4OH 
T.S.  shows  the  presence  of  chlorides. 

If  sulphides  or  cyanides  are  present,  add  an  excess  of  HNO3  T.S. 
to  a  small  portion  of  the  sample  and  boil,  under  a  hood,  to  expel 
H2S  or  HCN.  Then  add  AgNO3  T.S.  A  white  precipitate  shows 
the  presence  of  chlorides. 

If  bromides  or  iodides  are  present,  add  HNO3  T.S.  to  about  5  mils 
of  the  sample  and  boil  if  necessary  to  expel  H2S  or  HCN.  Add  an 
excess  of  AgNO3  T.S.  and  filter.  Pass  about  2  mils  of  NH4OH  T.S. 
once  through  the  precipitate  on  the  filter.  Any  silver  chloride 
quickly  dissolves  as  ammonio-silver  chloride.  Acidulate  the  filtrate 
with  HNO3.  A  white  precipitate  shows  the  presence  of  chlorides. 

I.  Acidulate  a  portion  of  the  sample,  add  HC1  T.S.,  and  add 
BaCl2  T.S.  A  white  precipitate  shows  the  presence  of  sulphates. 
If  a  precipitate  is  formed  upon  adding  HC1  T.S.,  it  should  be  filtered 
out  and  the  BaCl2  T.S.  should  be  added  to  the  clear  filtrate. 

J.  Add  H2SO4  T.S.  and  an  equal  volume  of  alcohol  to  a  portion 
of  the  sample  contained  in  a  porcelain  evaporating  dish  and  set  fire 
to  the  alcohol.  If  the  flame  is  colored  green,  borates  are  present. 

Barium  and  copper,  if  present,  must  be  removed  before  applying 
this  test  for  boric  acid. 

K.  Add  an  excess  of  H2SO4  T.S.  and  a  few  drops  of  alcohol  to 
a  portion  of  the  sample,  boil  and  observe  the  odor  of  the  vapor.  An 
odor  of  ethyl  acetate,  C^C^O^  shows  the  presence  of  acetates. 

L.  Mix  a  portion  of  the  sample  with  about  the  same  volume  of 
concentrated  H2SO4,  superimpose  FeSO4  T.S.  on  the  liquid  and  let 
stand.  A  brown  zone  between  the  two  liquids  shows  the  presence 
of  nitrates. 


SOLIDS. 

Solid  substances  must  generally  be  put  into  solution  before  they 
can  be  examined  for  bases  and  acids,  and  the  tests  must  be  applied 
to  the  solutions  obtained;  but  the  behavior  of  the  solid,  alone  or  with 
reagents,  will  often  give  information  leading  to  the  identification  of 
its  constituents,  and  various  methods  for  the  examination  of  solids 
have  been  devised.  The  method  given  here  is  aimed  at  bringing  the 
solid  into  solution  by  the  use  of  various  solvents.  Incidentally, 
some  information  of  analytical  value  may  be  obtained  by  observing 
the  directions  closely. 

SAMPLES  NUMBERS  19  AND  20. 
Qualitative  analysis  of  a  solid. 

The  sample  may  contain  any  of  the  bases  or  acids  considered  in  the 
course.  Interfering  organic  substances  are  absent. 

A.  The  sample  should  be  finely  powdered  with  a  mortar  and 
pestle,  or  otherwise,  if  it  is  not  already  in  a  powdered  condition. 

B.  Place  a  small  portion  of  the  powdered  sample  in  a  small  dry 
test-tube  and  heat  in  a  gas  flame,  carefully  observing  its  behavior 
and  the  color  and  odor  of  any  gas  or  vapor  given  off,  also  any  con- 
densation of  vapors  on  the  cool  portion  of  the  tube. 

If  no  apparent  change  occurs  in  the  solid  and  no  odor  is  observed, 
and  there  is  no  condensation  of  vapors,  then  organic  matter,  volatile 
bodies  and  water  of  crystallization  are  absent. 

If  the  solid  blackens  and  there  is  an  odor  of  something  burning, 
organic  matter  is  present.  If  organic  matter  were  present  it  would 
be  necessary  to  destroy  it,  by  methods  which  will  be  described  later, 
as  it  would  interfere  with  the  analytical  scheme. 

Compounds  of  some  of  the  metals  change  color  or  fuse  when 
heated,  but  the  changes  are  not  significant  enough  to  be  of  much 
analytical  value.  Such  metals,  if  present,  will  be  found  when  the 
groups  are  separated. 

If  the  substance  gives  off  a  vapor  which  condenses  on  the  cool  part 
of  the  tube,  volatile  substances  are  contained  in  the  solid,  or  are 
formed  upon  heating  it.  These  may  include  water;  any  compounds 
of  ammonium,  arsenic,  antimony  and  mercury;  or  some  compounds 
of  sulphur,  iron  and  iodine.  The  water  may  generally  be  recognized 
as  minute  drops  and  the  other  substances  will  be  found  in  their 
proper  places  when  the  scheme  of  analysis  is  applied. 


SOLIDS  107 

If  a  gas  is  evolved,  observe  its  odor  and  test  it  by  holding  a  splinter 
of  wood  or  a  cord  with  a  spark  at  the  end  in  it,  and  by  passing  some 
of  it  through  lime  water.  H2S,  SO2,  NH3,  NO2  and  some  other  gases 
may  be  recognized  by  their  odors.  If  the  gas  makes  a  spark  glow 
more  vividly  it  contains  more  oxygen  than  the  air  and  this  indicates 
the  presence  of  nitrates,  chlorates,  bromates,  iodates  or  peroxides. 
If  organic  matter  is  absent  and  the  gas  renders  lime  water  turbid, 
it  contains  carbon  dioxide  and  the  presence  of  carbonates  is  indicated. 

C.  Boil  about  1  gram  of  the  finely  divided  sample  with  about 
15  mils  of  water  in  a  beaker.    If  the  solid  does  not  completely 
dissolve  allow  the  residue  to  settle  and  decant  the  water  through  a 
filter.    Evaporate  a  few  drops  of  the  filtrate  to  dryness  on  a  watch- 
glass.     If  no  appreciable  deposit  is  left  on  the  watch-glass  the  sample 
is  insoluble  in  water.    In  such  a  case  proceed  with  paragraph  D, 
discarding  the  rest  of  the  filtrate.     If  any  appreciable  deposit  is  left 
on  the  watch-glass,  the  rest  of  the  filtrate,  or  the  solution  if  the 
sample  all  dissolved  in  water,  should  be  examined  for  the  bases  by 
the  tables  on  pages  82-85  and  for  acids  as  directed  on  pages  92,  97, 
99  and  103. 

D.  If  there  is  any.  residue  insoluble  in  water  from  C,  wash  it  with 
water  by  decantation  and  filtration  and  allow  to  drain,  discarding 
the  washings.    Transfer  two  small  portions  of  the  moist  powder  to 
test-tubes  and  boil  one  with  HC1  T.S.  and  the  other  with  HN03 
T.S.,  observing  which  of  the  acids  appears  to  be  the  better  solvent 
for  the  powder. 

If  there  is  an  effervescence  upon  adding  the  HC1  T.S.  to  the  powder, 
observe  the  color  and  odor  of  the  gas  and  pass  some  of  it  through  lime 
water.  A  greenish-yellow  gas  indicates  a  chlorate,  a  brown  gas 
indicates  a  nitrate,  and  a  purple  gas  indicates  an  iodide.  Gases 
with  characteristic  odors  are  yielded  by  salts  of  H2S,  H2SO3,  HBr, 
HI,  HCN  and  HC2H3O2.  If  the  gas  renders  lime  water  turbid, 
carbonates  are  indicated.  It  should  be  remembered  that  chlorine 
might  be  generated  from  the  HC1  used,  if  certain  oxidizing  agents 
are  present. 

Boil  the  rest  of  the  moist  powder  with  about  10  mils  of  the  acid 
which  appears  to  dissolve  it  the  better.  If  neither  acid  appears  to 
dissolve  any  of  the  powder,  or  both  dissolve  it  entirely,  use  HC1 
T.S.  in  preference  to  HNO3  T.S.  If  there  is  any  residue  insoluble 
in  the  acids,  allow  it  to  settle  and  decant  the  supernatant  liquid 
through  the  filter  used  for  the  water  solution  and  test  a  few  drops 
for  dissolved  matter  by  evaporating  on  a  watch-glass  as  before. 
If  any  appreciable  deposit  is  left  on  the  wratch-glass,  the  rest  of  the 
filtrate,  or  the  solution  if  all  the  residue  dissolved  in  the  acid,  should 
be  examined  for  bases  by  the  tables  on  pages  82-85,  and  for  acids  as 
directed  on  paged  92,  97,  99  and  103. 


108  SOLIDS 

E.  If  there  is  any  residue  insoluble  in  HC1  or  HNO3,  from  D,  it 
should  be  washed  by  decantation  and  filtration  as  before  and  digested 
with  about  5  mils  of  nitro-hydrochloric  acid.     Then  add  about  10 
mils  of  water  and  filter,  testing  the  filtrate  for  dissolved  matter  by 
evaporation  on  a  watch-glass  and  examine  for  bases  and  acids,  if 
any  appreciable  amount  of  the  solid  goes  into  solution. 

F.  If  sufficient  quantities  of  the  solvents  have,  been  used  and 
there  is  still  some  residue  left  from  E,  it  consists  of  insoluble  sulphates, 
chlorides  or  silicates.     To  decompose  these  transfer  the  washed 
and  drained  residue  to  a  porcelain  crucible,  mix  it  with  about  twice 
its  bulk  of  dry  Na2CO3  and  heat  to  fusion.     Then  allow  the  mass  to 
cool  and  boil  it  with  about  10  mils  of  water.     Filter  the  liquid  and 
examine  the  filtrate  for  acids  by  paragraph  G.     Dissolve  any  residue 
in  HC1  or  HNO3,  as  in  paragraph  D  and  examine  the  solution  for 
bases  by  the  tables  on  pages  82-85. 

G.  To  the  filtrate  from  F  add  an  excess  of  HNO3  T.S.  and  boil. 
If  no  precipitate  is  formed,  silicates  are  absent.     A  white  gelatinous 
precipitate  shows  the  presence  of  silicates.     If  a  precipitate  is  formed 
filter  the  liquid.     Divide  the  liquid  or  filtrate  into  two  portions. 

To  one  portion  add  BaCl2  T.S.  If  no  precipitate  is  formed,  sul- 
phates are  absent.  A  white  precipitate  shows  the  presence  of 
sulphates. 

To  the  other  portion,  add  AgN03  T.S.  If  no  precipitate  is 
formed,  chlorides  are  absent.  A  white  precipitate,  readily  soluble 
in  ammonia  water,  shows  the  presence  of  chlorides. 


QUALITATIVE  EXAMINATION  OF  OFFICIAL 
INOEGANIC  CHEMICALS. 

lin  speaking  of  the  purity  of  a  chemical  substance  we  refer  to  the 
kinds  and  amounts  of  other  substances  that  it  contains.  These 
other  substances  are  called  impurities. 

To  manufacture  an  absolutely  pure  chemical  compound,  it  would 
be  necessary  to  use  nothing  but  absolutely  pure  substances  and  not 
allow  the  compound  to  come  into  contact  with  anything  that  would 
contaminate  it  at  any  step  in  the  process  of  manufacture.  Such 
ideal  conditions  are  seldom  or  never  found,  as  natural  substances 
are  not  often  chemically  pure  and  most  chemicals  will  mix  with  more 
or  less  of  contaminating  substances  during  their  processes  of  manu- 
facture. Compounds  approximating  chemical  purity,  however,  are 
obtained  by  working  as  nearly  as  possible  under  the  ideal  conditions, 
or  by  purifying  compounds  of  a  lower  degree  of  purity.  In  general, 
water  is  considered  as  an  impurity  only  to  the  extent  that  it  affects 
the  strength  or  concentration  of  a  chemical. 

Many  grades  of  the  same  chemical  are  manufactured.  The  lowest 
commercial  grade  is  commonly  said  to  be  technically  pure  and  con- 
tains enough  of  the  chemical  to'  give  the  substance  identity,  but  may 
contain  considerable  amounts  of  impurities  Chemicals  of  this 
grade  are  made  in  large  quantities  and  are  used  in  many  kinds  of 
manufacturing  work.  The  highest  grade  of  a  chemical  is  generally 
labelled  chemically  pure  and  is  as  nearly  absolutely  pure  as  can  be 
obtained  under  manufacturing  conditions.  Such  chemicals  are 
relatively  high  in  price  and  are  only  used  when  such  a  high  grade  of 
purity  is  essential,  as  in  the  preparation  of  reagents  for  analytical 
work. 

For  medicinal  use,  chemicals  are  required  to  be  free  from  appre- 
ciable amounts  of  objectionable  impurities  and  of  proper  strength, 
but  the  United  States  Pharmacopoeia  and  the  National  Formulary 
do  not  require  that  they  be  of  chemically  pure  grade  in  many  cases, 
as  the  excessive  cost  of  such  chemicals  is  not  warranted,  provided 
that  the  impurities  are  not  harmful  or  excessive  in  amount.  A 
similar  rule  can  properly  be  applied  to  unofficial  chemicals.  The 
allowable  amount  of  impunties  is  stated  in  the  purity  rubric,  which  is 
a  part  of  the  definition  of  official  chemicals  and  states  the  minimum 
strength  required  for  concentrated  chemicals  and  the  minimum  and 
maximum  strengths  allowed  for  solutions  of  chemicals. 


110  EXAMINATION  OF  INORGANIC  CHEMICALS 

Qualitative  tests  on  official  chemicals  are  of  two  kinds,  namely, 
identity  tests  and  punty  tests,  these  being  used,  respectively,  to 
establish  the  identity  of  the  substance  under  examination,  and  to 
ascertain  the  presence  or  absence  of  excessive  amounts  of  objec- 
tionable impurities.  The  strengths  of  different  chemicals  are 
determined  by  assays,  which  belong  to  the  domain  of  quantitive 
analysis. 

The  directions  for  work  on  samples  21-25  are  taken  from  the 
United  States  Pharmacopoeia,  with  some  omissions  and  modifications, 
made  to  keep  the  work  within  the  scope  of  an  elementary  course  in 
qualitative  analysis. 

SAMPLE  NUMBER  21. 

Examination  of  granulated  ferrous  sulphate  for  identity  and  purity. 
(U.  S.  P.,  pp.  172  and  170.) 

The  tests  should  be  carried  out  as  directed,  the  following  report 
form  being  copied  and  filled  out  as  the  results  of  each  test  are 
obtained. 
Name Date 

Record  of  work  and  report  on  sample  No.  21. 

Qualitative  examination  of  granulated  ferrous  sulphate, 
(U.  S.  P.,  pages  172  and  170). 

Tests  for  identity: 
Color- 
Odor- 
Taste- 
Reaction  to  litmus — 
+  K3Fe(CN)6  T.S.- 
+  BaCl2  T.S.  and  HC1  T.S.- 

Tests  for  impurities: 
Heavy  metals — 
Free  acid — 

TESTS  FOR  IDENTITY. 

Grandulated  ferrous  sulphate  is  a  very  pale  bluish-green  crystal- 
line powder,  without  odor,  and  having  a  saline,  styptic  taste. 

An  aqueous  solution  of  the  salt  (1  in  20)  is  acid  to  litmus. 

An  aqueous  solution  of  the  salt,  even  when  highly  diluted,  gives 
with  potassium  ferricyanide  T.S.  a  blue  color  or  precipitate,  and  with 
barium  chloride  T.S.  a  white  precipitate  insoluble  in  hydrochloric 
acid. 


TESTS  FOR  IDENTITY  111 

TESTS  FOR  IMPURITIES. 

Dissolve  1  Gm.  of  the  salt  in  about  50  mils  of  distilled  water  con- 
taining 1  mil  of  diluted  sulphuric  acid,  heat  the  solution  to  boiling, 
oxidize  it  with  nitric  acid  and  then  mix  it  with  a  slight  excess  of 
ammonia  water  and  filter.  The  filtrate  is  colorless,  and,  after  acidu- 
lating with  hydrochloric  acid,  it  does  not  respond  to  the  test  for 
heavy  metals  (see  general  test  No.  3,  p.  118). 

Agitate  1  Gm.  of  ferrous  sulphate  in  small  fragments  during 
four  or  five  minutes,  with  10  mils  of  alcohol,  and  filter  the  mixture; 
the  filtrate  does  not  immediately  redden  moistened  blue  litmus 
paper  (free  acid) . 

SAMPLE  NUMBER  22. 

Examination  of  ammonium  bromide  for  identity  and  purity. 
(U.  S.  P.,  p.  43.) 

Tests  should  be  carried  out  as  described,  and  a  written  record  and 
report  made  as  the  work  progresses,  similar  to  that  directed  for 
sample  No.  21. 

TESTS  FOR  IDENTITY. 

Ammonium  bromide  occurs  in  colorless,  transparent,  prismatic 
crystals,  or  as  a  white,  crystalline  or  granular  powder;  odorless,  of  a 
pungent,  saline  taste. 

When  heated,  ammonium  bromide  volatilizes  without  fusing. 

An  aqueous  solution  of  the  salt  (1  in  20)  is  neutral  or  not  more  than 
slightly  acid  to  litmus. 

An  aqueous  solution  of  ammonium  bromide  when  gently  heated 
with  potassium  hydroxide  T.S.  evolves  ammonia. 

Silver  nitrate  T.S.  added  to  an  aqueous  solution  of  the  salt  (1  in 
10)  produces  a  yellowish-white  precipitate,  insoluble  in  nitric  acid 
or  in  a  moderate  excess  of  ammonia  water. 

TESTS  FOR  IMPURITIES. 

Add  1  mil  of  chloroform  to  10  mils  of  an  aqueous  solution  of  the 
salt  (1  in  20)  and  cautiously  introduce  chlorine  water,  which  has 
been  diluted  with  an  equal  volume  of  distilled  water,  drop  by  drop 
and  with  constant  agitation.  The  liberated  bromine  dissolves  in 
the  chloroform  and  imparts  to  it  a  yellow  or  orange  color,  which  is 
free  from  any  violet  tint  (iodide) . 

Drop  1  mil  of  diluted  sulphuric  acid  upon  about  1  Gm.  of  the 
powdered  salt;  no  yellow  color  appears  at  once  (bromate). 

A  blue  color  is  not  produced  at  once  on  adding  potassium  ferro- 
cyanide  T.S.  to  20  mils  of  an  aqueous  solution  of  the  salt  (1  in  250) 
(iron) . 


112  EXAMINATION  OF  INORGANIC  CHEMICALS 

Add  1  mil  of  potassium  sulphate  T.S.  to  10  mils  of  an  aqueous 
solution  of  ammonium  bromide  (1  in  20),  acidulated  with  acetic 
acid ;  no  turbidity  is  produced  immediately  (barium) ;  under  similar 
conditions  the  addition  of  barium  chloride  T.S.  produces  no  turbidity 
(sulphate) . 

An  aqueous  solution  of  the  salt  does  not  respond  to  the  test  for 
heavy  metals  (see  general  test  No.  3,  p.  118). 

SAMPLE  NUMBER  23. 

Examination  of  diluted  hydrochloric  acid  for  identity  and  purity. 
(U.  S.  P.,  p.  14.) 

The  tests  should  be  carried  out  as  described,  and  a  written  record 
and  report  made  as  the  work  progresses,  similar  to  that  directed 
for  sample  No.  21. 

TESTS  FOR  IDENTITY. 

Diluted  hydrochloric  acid  is  a  colorless,  odorless  liquid;  it  has  a 
strongly  acid  taste  and  is  strongly  acid  to  litmus. 

With  silver  nitrate  T.S.  it  yields  a  white,  curdy  precipitate, 
insoluble  in  nitric  acid,  but  readily  soluble  in  ammonia  water. 

TESTS  FOR  IMPURITIES. 

Add  1  mil  of  chloroform,  to  10  mils  of  diluted  hydrochloric  acid 
and  then  cautiously  introduce  chlorine  water,  which  has  been  diluted 
with  an  equal  volume  of  distilled  water,  a  drop  at  a  time,  with 
constant  agitation;  the  chloroform  remains  free  from  any  yellow, 
orange  or  violet  color  (bromide  or  iodide) . 

Add  1  mil  of  potassium  iodide  T.S.  and  1  mil  of  chloroform  to  10 
mils  of  diluted  hydrochloric  acid  and  agitate  the  mixture;  the 
chloroform  remains  free  from  any  violet  coloration  (free  chlorine 
or  bromine) . 

Two  mils  of  diluted  hydrochloric  acid,  diluted  with  seven  vol- 
umes of  distilled  water,  does  not  respond  to  the  test  for  heavy 
metals  (see  general  test  No.  3,  p.  118). 

Five  mils  of  diluted  hydrochloric  acid  diluted  with  12  mils  of 
distilled  water  (without  the  treatment  with  sulphuric  and  sulphur- 
ous acids)  meets  the  requirements  of  the  test  for  arsenic  (see  general 
test  No.  l,p.  115). 

Add  5  drops  of  barium  chloride  T.S.  to  a  mixture  of  3  mils  of 
diluted  hydrochloric  acid  and  5  mils  of  distilled  water;  neither 
turbidity  nor  precipitation  appears  within  one  hour  (sulphuric 
acid  or  sulphates).  At  the  end  of  this  period,  the  further  addition 
to  the  liquid  of  2  drops  of  tenth-normal  iodine  V.S.  produces  neither 
turbidity  nor  decoloration  of  the  iodine  (sulphurous  acid) . 


TESTS  FOR  IMPURITIES  113 

SAMPLE  NUMBER  24. 

Examination  of  bismuth  subnitrate  for  identity  and  purity. 
(U.  S.  P.,  pp.  82-83.) 

The  tests  should  be  carried  out  as  described  and  a  written  record 
and  report  made  as  the  work  progresses,  similar  to  that  directed  for 

sample  No.  21. 

TESTS  FOR  IDENTITY. 

Bismuth  subnitrate  is  a  white  powder;  odorless,  almost  tasteless. 

Bismuth  subnitrate  is  almost  ~  insoluble  in  water,  insoluble  in 
alcohol,  and  readily  dissolved  by  hydrochloric  or  nitric  acid. 

Its  solution  in  a  slight  excess  of  warm  nitrie  or  hydrochloric  acid 
produces  a  white  turbidity  when  added  to  25  volumes  of  distilled 
water. 

When  heated  to  redness  it  evolves  nitrous  vapors,  leaving  a  yellow 
residue,  which  is  blackened  by  hydrogen  sulphide  T.S. 

When  brought  in  contact  with  moistened  blue  litmus  paper,  the 
salt  shows  a  slightly  acid  reaction. 

TESTS  FOR  IMPURITIES. 

Boil  about  1  Gin.  of  bismuth  subnitrate  with  20  mils  of  a  mixture 
of  equal  parts  of  acetic  acid  and  distilled  water,  cool  the  solution 
and  filter.  Free  the  filtrate  from  bismuth  by  the  addition  of  hydro- 
gen sulphide,  boil  the  mixture  and  again  filter.  The  latter  filtrate 
leaves  no  appreciable  residue  -on  evaporation  and  gentle  ignition 
(alkalies  or  alkali  earths). 

Boil  about  0.1  Gm.  of  the  salt  with  5  mils  of  potassium  hydroxide 
T.S.;  no  odor  of  ammonia  is  perceptible  nor  does  the  vapor  turn 
moistened  red  litmus  paper  blue. 

The  residue  resulting  from  the  ignition  of  2  Gm.  of  the  salt  does 
not  respond  to  Bettendorf's  test  for  arsenic  (see  general  test  No.  2, 
p.  117). 

Add  3  Gm.  of  the  salt  to  3  mils  of  warm  nitric  acid;  no  efferves- 
cence occurs  (carbonate),  and  no  residue  remains  (insoluble  foreign 
salts).  Pour  this  solution  into  100  mils  of  distilled  water,  a  white 
precipitate  is  produced.  Filter,  evaporate  the  filtrate  on  a  water- 
bath  to  30  mils,  again  filter  the  liquid  and  divide  the  new  filtrate 
into  portions  of  5  mils  each. 

Mix  one  portion  with  an  equal  volume  of  diluted  sulphuric  acid; 
it  does  not  become  cloudy  (lead) . 

Precipitate  another  portion  with  a  slight  excess  of  ammonia  water; 
the  supernatant  liquid  does  not  exhibit  a  bluish  tint  (copper). 

Another  portion  is  not  immediately  affected  by  barium  nitrate 
T.S.  (sulphate).    With  hydrochloric  acid  no  precipitate  is  formed 
8 


114  EXAMINATION  OF  INORGANIC  CHEMICALS 

which  is  insoluble  in  a  slight  excess  of  the  latter,  but  soluble  in 
ammonia  water  (silver) . 

SAMPLE  NUMBER  25. 

Examination  of  alum  for  identity  and  purity. 
(U.S.  P.,  pp.  39  40). 

The  tests  should  be  carried  out  as  described  and  a  written  record 
and  report  made  as  the  work  progresses,  similar  to  that  directed 
for  sample  No.  21. 

TESTS  FOR  IDENTITY. 

Ammonium  alum  and  potassium  alum  both  occur  in  large,  color- 
less crystals,  crystalline  fragments,  or  as  white  powders;  alum  is 
odorless  and  has  a  sweetish  and  strongly  astringent  taste. 

Ammonium  alum. — Potassium  hydroxide  T.S.  added  to  an  aqueous 
solution  of  ammonium  alum  (1  in  20)  at  first  causes  a  precipitate, 
which  completely  dissolves  in  an  excess  of  the  reagent,  ammonia 
being  evolved. 

Potassium  alum. — Potassium  alum  imparts  a  violet  color  to  a 
non-luminous  flame. 

The  addition  of  sodium  bitartrate  T.S.  to  a  saturated  solution  of 
potassium  alum  produces,  sometimes  slowly,  a  white  crystalline 
precipitate. 

Potassium  hydroxide  T.S.  added  to  an  aqueous  solution  of  potas- 
sium alum  (1  in  20)  at  first  causes  a  precipitate,  which  completely 
dissolves  in  an  excess  of  the  reagent,  but  no  ammonia  is  evolved. 

Both  ammonium  alum  and  potassium  alum  conform  to  the  following 
tests: 

An  aqueous  solution  of  alum  (1  in  20)  is  acid  to  litmus. 

An  aqueous  solution  of  alum  (1  in  20)  yields  with  ammonia  water 
a  white  gelatinous  precipitate  almost  insoluble  in  an  excess  of 
ammonia  water. 

With  barium  chloride  T.S.,  an  aqueous  solution  of  alum  (1  in  20) 
yields  a  white  precipitate  insoluble  in  hydrochloric  acid. 

TESTS  FOR  IMPURITIES. 

An  aqueous  solution  of  alum  does  not  respond  to  the  test  for 
heavy  metals  (see  general  test  No.  3,  p.  118).  An  aqueous  solution 
of  alum  meets  the  requirements  of  the  test  for  arsenic  (see  general 
test  No.  l,p.  116). 

Add  5  drops  of  potassium  ferrocyanide  T.S.  to  20  mils  of  an 
aqueous  solution  of  alum  (1  in  150);  no  blue  color  is  produced  at 
once  (iron). 


GENERAL  QUALITATIVE  TESTS  OF  THE 
UNITED  STATES  PHAEMACOPCEIA. 

1.  Arsenic  Test.  —  Test  Apparatus. — Prepare  several  generators, 
equipped  with  tubes,  etc.,  as  described  below.  Select  as  a  generator 
a  bottle  of  about  50  mils  capacity,  having  a  mouth  about  2.5  cm.  in 
diameter  and  provide  a  well-fitting  rubber  stopper,  suitably  per- 
forated. In  the  perforation  in  this  stopper  insert  a  vertical  exit 
tube  about  13  cm.  in  total  length  and  1  cm.  in  diameter  throughout 
the  upper  portion  (for  about  10  cm.)  and  constricted  at  its  lower 
extremity  to  a  tube  of  about  3  cm.  in  length  and  about  5  mm.  in 
diameter.  This  latter  tube  should  extend  but  slightly  below  the 
stopper.  In  the  lower  part  of  the  wider  exit  tube  insert  a  small 
pledget  of  dry  glass  wool  and  then  a  strip  of  the  freshly  prepared  but 
dry  lead  acetate  test-paper  rolled  into  a  coil,  and  above  this  a  plug 
of  the  moist  (not  wet)  lead  acetate  glass  wool.  In  the  upper 
extremity  of  this  tube  insert  through  a  perforated  cork  stopper,  a 
glass  tube  12  cm.  in  length,  having  an  internal  diameter  of  about 
3  mm.  Place  the  mercuric  bromide  test-paper  in  this  tube,  bending 
or  creasing  the  upper  portion  of  the  strip  so  that  it  will  retain  its 
position.  The  strip  should  extend  within  about  2  cm.  of  the  per- 
forated cork  stopper  and  must  not  be  introduced  into  the  tube  until 
ready  to  start  the  test.  This  tube  should  be  thoroughly  cleaned 
and  dried  each  time  it  is  used. 

Standard  Arsenic  Test-solution.  —  Dissolve  0.1  Gm.  of  arsenic 
trioxide,  which  has  been  finely  pulverized,  dried  in  a  desiccator, 
and  accurately  weighed,  in  about  5  mils  of  a  20  per  cent,  solution  of 
sodium  hydroxide  (free  from  arsenic) .  Neutralize  the  solution  with 
diluted  sulphuric  acid  (free  from  arsenic),  add  10  mils  of  the  same 
acid  and  sufficient  recently  boiled  distilled  water  to  bring  the 
volume  of  the  solution  to  exactly  1000  mils  at  25°  C.  Accurately 
measure  10  mils  of  this  solution,  transfer  it  to  a  1000  mil  flask, 
add  10  mils  of  diluted  sulphuric  acid  (free  from  arsenic)  and  make 
up  the  volume  with  recently  boiled  distilled  water  to  exactly  1000 
mils  at  25°  C.  Employ  this  solution,  containing  0.001  mg.  of  arsenic 
trioxide  in  each  1  mil  (at  25°  C.),  in  preparing  the  standard  stain. 
Keep  this  solution  in  a  glass-stoppered  bottle.  It  is  advisable  to 
make  fresh  solutions  whenever  new  standard  stains  are  to  be 
prepared, 


116  TESTS  OF   UNITED  STATES  PHARMACOPCEIA 

Preparation  of  the  Chemical  to  be  Tested. — Add  1  mil  of  a  mixture 
of  equal  volumes  of  concentrated  sulphuric  acid  and  distilled  water, 
to  5  mils  of  an  aqueous  solution  of  the  chemical  (1  in  25)  or  to  a 
solution  in  5  mils  of  distilled  water  of  the  residue  remaining  after 
any  special  treatment  that  may  be  directed.  This  acidulation  is  not 
necessary  in  the  case  of  inorganic  acids.  Now,  unless  specially 
directed,  add  10  mils  of  a  saturated  aqueous  solution  of  sulphurous 
acid.  Evaporate  this  liquid  in  a  small  beaker,  on  a  water-bath, 
until  it  is  free  from  sulphurous  acid  and  has  been  reduced  to  about 
2  mils  in  volume.  Dilute  this  evaporated  liquid  to  about  5  mils 
with  distilled  water. 

THE  TEST  FOR  ARSENIC. — Preparation  of  the  Standard  Stain. — 
Introduce  into  the  generator  from  8  to  10  Gm.  of  zinc  followed  by 
25  mils  of  dilute  sulphuric  acid,  prepared  by  mixing  1  part  of  con- 
centrated sulphuric  acid,  free  from  arsenic,  with  4  parts  of  dis- 
tilled water,  and  5  drops  of  acid  stannous  chloride  T.S.  Add 
at  once  2  mils  (accurately  measured  at  25°  C.)  of  the  standard 
arsenic  T.S.  and  immediately  insert  the  stopper  containing  the 
exit  tube,  into  which  have  been  placed  the  glass-wool  pledget, 
the  dry  lead  acetate  test-paper,  the  moist  lead  acetate  glass  wool, 
and  the  mercuric  bromide  test-paper,  as  described  under  the  Test 
Apparatus.  Should  the  evolution  of  the  gas  be  violent  at  first, 
check  the  reaction  by  immersing  the  bottle  in  cold  water.  Should 
the  reaction  subside,  increase  it  by  placing  the  bottle  in  warm  water. 
If  the  reaction  be  too  violent,  the  stain  will  spread  and  not  form  a 
distinctive  band,  thus  making  the  comparisons  of  color  intensity 
difficult.  After  the  test  has  continued  for  forty-five  minutes, 
remove  the  mercuric  bromide  test-paper  and  place  it  in  a  clean, 
dry  tube  for  comparison.  This  stain  represents  0.002  mg.  of 
arsenic  trioxide  in  addition  to  any  stain  produced  by  the  reagents. 
The  stain  from  the  reagents  should  scarcely  be  perceptible  when 
determined  by  a  blank  experiment.  Since  light,  heat  and  moisture 
cause  the  stain  to  fade  rapidly,  comparison  should  be  made  as  soon 
as  possible.  The  stained  test-papers  may  be  preserved  by  either 
dipping  in  hot,  melted  paraffin  or  keeping  them  over  phosphoric 
anhydride  protected  from  light. 

Testing  the  Chemical. — Introduce  into  another  generator  from  8 
to  10  Gm.  of  the  zinc,  followed  by  25  mils  of  the  dilute  sulphuric 
acid  (1  to  4)  and  5  drops  of  acid  stannous  chloride  T.S.  Insert  the 
stopper  containing  the  exit  tube  charged  with  the  test-papers  and 
glass  wool,  as  just  described.  Then  add  5  mils  of  the  solution 
to  be  tested,  previously  treated  as  directed,  under  Preparation  of  the 
Chemical,  and  immediately  insert  the  stopper  charged  with  the  test- 
papers  and  glass  w^ool.  When  the  evolution  of  hydrogen  has  pro- 


ARSENIC  TEST  117 

ceeded  actively  for  forty-five  minutes,  remove  the  mercuric  bromide 
test-paper  and  carefully  compare  it  with  the  standard  stain  prepared 
as  described  above.  The  stain  produced  by  the  chemicals  tested 
does  not  exceed  in  length  or  intensity  of  color  that  prepared  as  the 
standard,  indicating  not  more  than  1  part  of  arsenic  in  100,000 
parts  of  the  substance  tested. 

Interfering  Chemicals. — Antimony,  if  present  in  the  substance 
tested,  will  produce  a  gray  stain.  Sulphites,  sulphides,  thiosulphates , 
and  other  compounds  which  liberate  hydrogen  sulphide  or  sul- 
phurous acid  when  treated  with  sulphuric  acid,  must  be  oxidized 
by  means  of  nitric  acid  and  then  reduced  by  means  of  sulphurous 
acid  as  directed  under  Preparation  of  the  chemical  to  be  tested, 
before  introducing  them  into  the  apparatus.  Sulphur  compounds 
as  well  as  hydrogen  phosphide  give  a  bright  yellow  band  on  the  test 
paper.  If  sulphur  compounds  are  present,  a  simultaneous  darkening 
of  the  lead  acetate  test  paper  and  glass  wool  will  occur.  If  such  is 
the  case,  the  operation  as  directed  under  Preparation  of  the  chemical 
to  be  tested  must  be  repeated  upon  a  fresh  portion  of  the  solution 
being  tested,  using  greater  care  in  effecting  the  complete  "emoval  of 
the  sulphurous  acid.  In  testing  hypophosphites  special  care  should 
be  observed  to  completely  oxidize  the  solution  being  tested  as 
directed,  otherwise  a  yellow  stain  may  be  produced  through  the 
evolution  of  hydrogen  phosphide,  which  might  be  confused  with  the 
orange-yellow  color  produced  by  arsenic.  Compounds  containing 
antimony  should  be  tested  for  arsenic  by  Bettendorf's  Test  (see  Test 
No.  2). 

The  test  apparatus  should  be  thoroughly  cleaned  and  dried 
immediately  before  and  after  use. 

2  Arsenic  test,  Bettendorf's.— This  test  is  employed  only  in  test- 
ing salts  of  bismuth  and  compounds  containing  antimony  for  the 
presence  of  arsenic. 

To  a  solution  of  the  prescribed  quantity  of  the  substance  to  be 
tested  in  5  mils  of  concentrated  hydrochloric  acid  contained  in  a 
clean  test  tube,  add  10  mils  of  saturated  stannous  chloride  T.S. 
which  has  been  freshly  prepared,  and  set  it  aside  for  thirty  minutes. 
If  arsenic  is  present  in  non-permissible  amounts,  a  brownish  tint  or 
precipitate  will  be  seen  when  the  tube  is  placed  over  a  white  surface 
and  the  solution  viewed  from  above,  comparison  being  made  with 
a  mixture  of  5  mils  of  concentrated  hydrochloric  acid  and  10  mils 
of  concentrated  stannous  chloride  T.S.,  prepared  as  directed  above. 

NOTE  :  It  is  absolutely  necessary  that  for  this  test  the  solution  of 
stannous  chloride  be  freshly  prepared,  and  that  nitrates,  sulphates, 
sulphites,  sulphides  and  compounds  of  mercury,  gold,  and  selenium 
be  absent  from  the  reagents,  and  from  the  chemicals  being  tested. 


118  TESTS  OF  UNITED  STATES  PHARMACOPEIA 

3  Test  for  heavy  metals.— Chemicals.— This  test  is  to  be  used  to 
detect  the  presence  of  undesirable  metallic  impurities  in  official 
chemical  substances  or  their  solutions;  these  must  not  respond 
affirmatively  within  the  stated  time.  Acidulate  10  mils  of  a  solution 
of  the  substance  in  distilled  water  (1  in  50),  contained  in  a  test  tube 
of  about  40  mils  capacity  and  of  about  2.5  cm.  diameter,  with  1  mil 
of  diluted  hydrochloric  acid  (unless  otherwise  directed),  warm  it 
to  about  50°  C.,  add  an  equal  volume  of  freshly  prepared  hydrogen 
sulphide  T.S .,  stopper,  and  allow  the  mixture  to  stand  at  35°  C. 
for  half  an  hour.  At  the  end  of  this  time  the  mixture  should  still 
possess  the  odor  of  hydrogen  sulphide;  if  not,  it  should  be  thoroughly 
saturated  with  the  gas  and  again  set  aside  for  half  an  hour.  The 
color  produced,  if  any,  is  not  greater  than  that  observed  by  a  blank 
test  made  in  the  same  manner  and  with  the  same  quantities  of  the 
reagents  (omitting  the  solution  to  be  tested);  the  solutions  being 
viewed  crosswise  by  reflected  light  while  held  against  a  white  sur- 
face. A  slight  turbidity  due  to  separation  of  sulphur  from  the 
hydrogen  sulphide  may  occur. 


DESTEUCTION  OF  ORGANIC  MATTER. 

If  organic  matter  is  present  in  a  sample,  it  is  necessary  to  destroy 
it  before  proceeding  with  a  qualitative  analysis  for  inorganic  bases 
and  acids.  There  are  several  methods  for  accomplishing  this,  of 
which  the  most  generally  useful  are  the  following : 

A.  Mix  the  sample  with  an  excess  of  concentrated  HC1,  add  some 
crystals  of  KC1O3  and  boil  under  a  hood  until  the  mixture  is  decolor- 
ized or  a  portion  does  not  carbonize  upon  drying  and  igniting.    Add 
additional  HC1  and  KC1O3  as  necessary,  avoiding  a  large  excess. 
Then  evaporate  the  liquid  to  dryness  on  a  water  bath  and  examine 
the  residue  by  the  method  for  the  qualitative  analysis  of  a  solid. 

This  method  destroys  the  organic  matter,  or  changes  it  so  that  it 
does  not  interfere  with  the  analysis.  The  metals  present  are 
oxidized  to  their  highest  valences,  and  ammonium  salts  are  volati- 
lized. Allowance  must  also  be  made  for  the  potassium  and  chlorine 
compounds  added. 

B.  Mix  tbe  sample  with  about  4  times  its  weight  of  concentrated 
H2S04  and  twice  its  weight  of  concentrated  HNO3,  digest  at  the 
ordinary  temperature  for  several  hours,  and  then  heat  in  a  hood  until 
the  white  fumes  of  H2SO4  are  given  off.     If  the  mixture  still  contains 
black  carbonized  organic  matter,  add  more  HNO3  and  repeat  the 
heating  until  white  fumes  are  given  off,  repeating  this  treatment 
with  additional  HNO3  and  heating  several  times  if  necessary.    When 
the  mixture  ceases  to  show  the  presence  of  carbon,  cautiously  add 
about  an  equal  volume  of  water  and  concentrate  until  the  white 
fumes  of  H2SO4  are  given  off,  then  cautiously  add  four  or  five  times 
the  volume  of  water  and  filter,  if  necessary.    Examine  the  filtrate 
by  the  method  for  the  qualitative  analysis  of  a  liquid.    Any  residue, 
on  the  filter  paper,  may  contain  silica  or  the  sulphates  of  lead, 
barium,  strontium  and  calcium.    Examine  it  by  the  method  for 
the  qualitative  analysis  of  a  solid,  beginning  with  paragraph  F  on 
page  108. 

This  is  the  most  dependable  method  for  the  destruction  of  organic 
matter.  Allowance  must  be  made,  however,  for  the  H2SO4  and 
HN03  used,  and  it  is  often  necessary  to  use  both  of  these  methods, 
or  others,  in  making  a  careful  qualitative  analysis  of  a  substance 
containing  organic  matter. 


REAGENTS  AND  TEST  SOLUTIONS. 

On-  the  following  pages  will  be  found  directions  for  preparing  the 
test  solutions  called  for  in  this  book.  Most  of  these  are  of  the 
same  strength  as  is  specified  in  the  United  States  Pharmacopoeia, 
but  for  performing  the  Pharmacopceial  tests  on  official  chemicals, 
the  reagents  should  be  made  according  to  the  directions  of  the 
Pharmacopoeia  IX,  pages  521-583. 

In  preparing  reagents  and  test  solutions,  only  pure  chemicals 
and  distilled  water  should  be  used. 

Acetic  acid,  HC2H3O2. — Mix  official  acetic  acid,  containing  about 
36  per  cent,  of  absolute  HC2H3O2  with  twice  its  volume  of  distilled 
water. 

Ammonium  hydroxide  test  solution,  NH4OH  or  NH3  +  H2O.— Use 
the  official  ammonia  water,  containing  10  per  cent,  of  NH3. 

Ammonium  carbonate  test  solution,  (NH4)2CO3.— Dissolve  20  Gm. 
of  official  ammonium  carbonate  in  a  mixture  of  20  mils  of  ammonia 
water  and  65  mils  of  distilled  water,  and  add  sufficient  distilled  water 
to  measure  100  mils. 

Ammonium  chloride  test  solution,  NH4C1.— Dissolve  10  Gm.  of 
ammonium  chloride  in  sufficient  distilled  water  to  make  100  mils. 

Ammonium  molybdate  test  solution,  (NH4)6Mo7O2  +  4H2O.— Mix 
6.5  Gm.  of  finely  powdered  molybdic  acid  with  14  mils  of  distilled 
water  and  14.5  mils  of  stronger  ammonia  water  to  effect  solution. 
Cool  and  slowly  add  the  solution,  in  small  portions  with  agitation, 
to  a  well  cooled  mixture  of  32  mils  of  nitric  acid  and  40  mils  of  dis- 
tilled water.  Allow  the  solution  to  stand  for  twenty-four  hours 
and  then  filter  through  asbestos. 

Preserve  the  test  solution  in  the  dark,  and,  if  a  sediment  should 
form  in  it  after  some  days,  carefully  decant  the  clear  solution. 
This  solution  should  be  tested  at  frequent  intervals.  Add  2  mils  of 
sodium  phosphate  T.S.  to  5  mils  of  the  reagent;  an  abundant 
yellow  precipitate  forms  either  at  once  or  upon  slight  warming.  If 
only  a  slight  precipitation  or  yellow  opalescence  results,  the  reagent 
must  be  rejected.  When  employed  as  a  reagent,  ammonium 
molybdate  T.S.  is  always  added  in  large  excess  to  the  solution  being 
tested,  the  latter  having  previously  been  strongly  acidified  with 
nitric  acid. 

Ammonium  oxalate  test  solution  (NH4)2C2O4  +  H2O.— Dissolve 
4  Gm.  of  ammonium  oxalate  in  sufficient  distilled  water  to  measure 
100  mils. 


FERROUS  SULPHATE  TEST  SOLUTION  121 

Ammonium  polysulphide  test  solution,  Yellow  Ammonium  Sul- 
phide, (XH4)2Sx. — Saturate  three  parts  of  ammonia  water  with 
hydrogen  sulphide,  forming  a  solution  of  ammonium  hydrogen 
sulphide,  XH4HS.  Add  to  the  solution  two  additional  parts  of 
ammonia  water,  forming  a  solution  of  normal  or  colorless  ammonium 
sulphide,  (NH4)2S.  Finally  add  a  small  quantity  of  precipitated 
sulphur  and  preserve  the  solution  in  small,  amber,  glass-stoppered 
bottles. 

Barium  chloride  test  solution,  BaCl2  +  2H2O. — Dissolve  10  Gm. 
of  barium  chloride  in  sufficient  distilled  water  to  make  100  mils. 

Barium  hydroxide  test  solution,  Ba(OH)2  +  8H2O. — A  freshly  made 
saturated  solution  of  barium  hydroxide  in  distilled  water. 

Barium  nitrate  test  solution,  Ba(NO3)2. — Dissolve  5  Gm.  of  barium 
nitrate  in  sufficient  distilled  water  to  make  100  mils. 

Bromine  test  solution  (bromine  water),  Br. — A  saturated  solution  of 
bromine,  prepared  by  adding  3  mils  of  bromine  to  100  mils  of  cold 
water.  The  solution  should  be  stored  in  a  glass-stoppered  bottle 
in  a  cool  place,  protected  from  light  and  should  be  shaken  well, 
allowing  the  excess  of  bromine  to  settle  before  using. 

Calcium  chloride  test  solution,  CaCl2  +  2H2O.— Dissolve  10  Gm. 
of  calcium  chloride  in  sufficient  distilled  water  to  measure  100  mils. 

Calcium  hydroxide  test  solution  (lime  water),  Ca(OH)2. — A  nearly 
saturated  aqueous  solution  of  calcium  hydroxide. 

Calcium  sulphate  test  solution,  CaSO4  +  2H2O. — Mix  5  mils  of 
diluted  sulphuric  acid  with  200  mils  of  distilled  water,  add  1  Gm.  of 
calcium  carbonate  and  shake:  when  the  acid  is  completely  neutral- 
ized filter  the  mixture. 

Chlorine  test  solution  (chlorine  water),  Cl. — A  saturated  aqueous 
solution  of  chlorine,  prepared  by  generating  the  gas  with  manganese 
and  hydrochloric  acid,  washing  and  conducting  it  into  cold  water 
until  saturated.  The  solution  should  be  preserved  in  small  amber 
glass-stoppered  bottles  in  a  cool,  dark  place  and  it  should  have  a 
strong  odor  of  chlorine  when  used* 

Cobaltous  chloride  test  solution,  Co4Cl  +  6H20. — Dissolve  2  Gm. 
of  cobaltous  chloride  with  the  aid  of  1  mil  of  hydrochloric  acid  in 
sufficient  distilled  water  to  measure  100  mils. 

Copper  sulphate  test  solution,  CuSO4  +  5H2O. — Dissolve  10  Gm. 
of  copper  sulphate  in  sufficient  distilled  water  to  make  100  mils. 

Ferric  chloride  test  solution,  FeCl3  +  6H2O. — Dissolve  10  Gm.  of 
ferric  chloride  in  sufficient  water  to  make  100  mils. 

Ferrous  sulphate  test  solution.,  FeSO4  +  7H2O.— Dissolve  10  Gm. 
of  clear  crystals  of  ferrous  sulphate  in  sufficient  distilled  water  to 
make  100  mils.  This  solution  is  to  be  freshly  prepared  immediately 
before  use. 


122  REAGENTS  AND  TEST  SOLUTIONS 

Hydrochloric  acid  test  solution,  HC1.— Use  the  official  diluted 
hydrochloric  acid  of  10  per  cent,  strength. 

Hydrogen  sulphide  test  solution,  H2S.— A  saturated,  aqueous  solu- 
tion of  hydrogen  sulphide.  To  prepare  about  1000  mils  of  the  solution, 
treat  20  Gm.  of  ferrous  sulphide  in  a  suitable  apparatus  with  a  mixture 
of  20  mils  of  sulphuric  acid  and  250  mils  of  distilled  water,  pass  the  gas 
through  a  drying  tube  filled  with  granulated  calcium  chloride,  then 
from  this  through  a  tube  of  about  8  millimeters  diameter  and  40  centi- 
meters in  length,  which  contains  about  5  Gm.  of  coarsely  pulverized 
iodine  mixed  with  glass  wool,  and  finally  through  a  wash  bottle  which 
contains  a  small  quantity  of  potassium  iodide  T.S.  The  gas  thus  puri- 
fied is  conducted  nearly  to  the  bottom  of  a  bottle  of  the  capacity  of 
about  1500  mils  containing  1000  mils  of  cold  distilled  water.  Shake 
the  bottle  occasionally  to  facilitate  the  solution  of  the  gas.  When 
the  gas  is  no  longer  absorbed,  transfer  the  solution  to  small,  dark, 
amber-colored  bottles,  fill  nearly  to  the  top ;  pass  a  stream  of  purified 
hydrogen  sulphide  for  a  few  minutes  through  each,  and  then  at  once 
stopper  them  tightly,  and  preserve  them  afterwards  in  a  cool  and 
dark  place.  Do  not  use  this  solution  unless  it  retains  a  strong  odor 
of  hydrogen  sulphide,  and,  when  added  to  an  equal  volume  of  ferric 
chloride  T.S.,  produces  at  once  a  copious  precipitate  of  sulphur. 

Iodine  test  solution,  I.— Dissolve  1  Gm.  of  iodine  and  3  Gm.  of 
potassium  iodide,  in  50  mils  of  distilled  water. 

Lead  acetate  glass  wool,  Pb(C2H302)2  +  3H2O. — Immerse  glass 
wool  in  a  mixture  of  equal  parts  of  lead  acetate  T.S.  and  water  and 
remove  the  excess  of  liquid  by  pressing  it  between  filter  paper.  It 
should  be  prepared  immediately  before  it  is  to  be  used. 

Lead  acetate  test  paper,  Pb(C2H3O2)2  +  3H2O. — Immerse  strips 
of  heavy  white  filter  paper,  6  cm.  in  width  and  8  cm.  in  length,  in  a 
mixture  of  equal  parts  of  lead  acetate  T.S.  and  distilled  water,  drain 
off  the  excess  of  liquid  and  dry  the  paper  in  an  oven  at  100°  C., 
avoiding  contact  with  metal. 

Lead  acetate  test  solution,  Pb(C2H3O2)2  +  3H2O.— Dissolve  10 
Gm.  of  clear  transparent  crystals  of  lead  acetate,  free  from  adhering 
lead  carbonate,  in  sufficient  distilled  water  to  make  100  mils.  Pre- 
serve the  solution  in  well-stoppered  bottles. 

Litmus  paper  and  test  solution.— Exjiaust  powdered  litmus  with 
three  separate  and  successive  portions  (each  equal  to  about  4  times 
its  weight)  of  boiling  alcohol  (which  removes  the  undesirable  color 
erythrolitmin),  each  extraction  lasting  for  about  one  hour.  After 
draining  off  the  alcohol,  digest  the  residue  with  about  an  equal 
weight  of  cold  water  and  filter.  (This  blue  solution,  which  contains 
some  alkali,  after  being  acidulated,  may  be  used  to  make  red  litmus 
paper.)  Finally,  extract  the  residue  with  about  5  times  its  weight 
of  boiling  distilled  water,  and,  after  thoroughly  cooling,  filter.  The 


MERCURIC  CHLORIDE  TEST  SOLUTION  123 

addition  of  1  drop  of  hundredth-normal  acid  or  alkali  V.S.  to  50 
mils  of  distilled  water  containing  5  drops  of  the  indicator  produces 
a  distinct  change  in  color.  Preserve  the  filtrate,  as  a  test  solution, 
in  wide-mouthed  bottles  stoppered  with  loose  plugs  of  purified 
cotton  so  as  to  exclude  dust  but  admit  air.  The  latter  must  be  free 
from  acid  or  ammoniacal  vapors.  The  blue  color  of  litmus  test 
solution  is  changed  by  acids  to  red,  and  this  red  color  by  the  addition 
of  alkalies  is  restored  to  blue. 

Litmus  Paper,  Blue. — Impregnate  with  the  test  solution  just 
described  strips  of  white  filter  paper,  and  dry  them  by  suspending 
them  on  lines  of  clean  twine,  in  an  atmosphere  free  from  acid  or 
ammoniacal  vapors.  This  paper  must  quickly  respond  to  a  two 
hundred  and  fiftieth-normal  acid  V.S. 

Litmus  Paper,  Red— Prepare  this  with  the  same  kind  of  paper  and 
in  the  manner  described  under  Litmus  Paper,  Blue,  having  added  to 
the  test  solution  used  to  impregnate  the  paper  just  sufficient  of  a 
highly  diluted  solution  of  hydrochloric  acid  to  impart  to  it  a  faint 
red  tint.  Neither  blue  nor  red  litmus  paper  should  have  an  intense 
color.  Preserve  the  test  paper  in  bottles,  so  as  to  exclude  dust  and 
acid  or  ammoniacal  vapors.  This  paper  must  quickly  respond  to  a 
two  hundred  and  fiftieth-normal  alkali  V.S. 

Magnesia  mixture,  MgCl2  +  6H2O,  NH4C1,  NH4OH.- Dissolve 
5.5.  Gm.  of  magnesium  chloride  and  7  Gm.  of  ammonium  chloride 
in  65  mils  of  distilled  water,  add  35  mils  of  ammonia  water,  set  the 
mixture  aside  for  a  few  day&  in  a  well-stoppered  vessel,  and  filter. 
If  not  perfectly  clear,  filter  the  solution  before  using. 

Magnesium  sulphate  test  solution,  MgSCX  +  7H20. — Dissolve 
10  Gm.  of  magnesium  sulphate  in  sufficient  distilled  water  to 
measure  100  mils. 

Manganese  sulphate  test  solution,  MnSCX  +  4H2O. — Dissolve  10 
Gm.  of  manganese  sulphate  in  50  mils  of  distilled  water  and  add 
sufficient  diluted  sulphuric  acid  to  make  100  mils. 

Mercuric  bromide  test  paper,  HgBr2. — Cut  stiff,  heavy  quanti- 
tative filter  paper  into  strips  3  mm.  in  width  and  about  12  cm.  in 
length.  Immerse  these  strips  for  five  minutes  in  alcoholic  mercuric 
bromide  T.S.  Remove  the  excess  of  solution  by  pressing  the  strips 
between  filter  paper  and  then  dry  them  quickly  on  glass  in  an  oven 
heated  to  100°  C.  Place  the  strips  at  once  in  a  wide  mouthed 
bottle  and  stopper  it  securely. 

Mercuric  bromide  test  solution,  alcoholic,  HgBr2. — Dissolve  5  Gm. 
of  mercuric  bromide  in  100  mils  of  alcohol,  employing  a  gentle  heat 
to  facilitate  solution.  Keep  it  in  glass  stoppered  bottles  protected 
from  the  light. 

Mercuric  chloride  test  solution,  HgCl2. — Dissolve  5  Gm.  of  mercuric 
chloride  in  sufficient  distilled  water  to  measure  100  mils. 


124  REAGENTS  AND  TEST  SOLUTIONS 

Mercuric  potassium  iodide  test  solution  (Mayer's  reagent),  HgI2  + 
2KI. — Dissolve  1.358  Gm.  of  mercuric  chloride  in  60  mils  of  dis- 
tilled water,  and  5  Gm.  of  potassium  iodide  in  10  mils  of  distilled 
water.  Mix  the  two  solutions,  arid  then  add  sufficient  distilled 
water  to  measure  100  mils. 

Mercuric  potassium  iodide  test  solution,  alkaline  (Nessler's  Reagent). 
— Dissolve  10  Gm.  of  potassium  iodide  in  10  mils  of  distilled  water, 
and  add  gradually  in  portions  a  saturated  aqueous  solution  of 
corrosive  mercuric  chloride  with  constant  agitation,  until  a  slight  red 
precipitate  remains  undissolved;  to  this  mixture  add  30  Gm.  of  potas- 
sium hydroxide  and,  when  solution  has  taken  place,  1  mil  more  of 
the  saturated  aqueous  solution  of  mercuric  chloride.  Dilute  this 
solution  with  distilled  water  until  it  measures  200  mils.  Allow  the 
precipitate  to  subside,  and  draw  off  the  clear  fluid.  Two  mils  of 
this  reagent,  when  added  to  50  mils  of  distilled  water  containing 
0.05  mg.  of  ammonia,  produces  at  once  a  yellowish  brown  coloration. 

Mercurous  nitrate  test  solution,  HgNO3. — Mix  10  Gm.  of  pure 
mercury  with  5  mils  of  nitric  acid  and  5  mils  of  distilled  water,  in  a 
porcelain  evaporating  dish  and  set  it  aside  for  twenty-four  hours  in  a 
cool,  dark  room.  Separate  and  drain  the  crystals  of  mercurous 
nitrate,  and  dissolve  them  in  100  mils  of  distilled  water.  Preserve 
the  solution  in  a  dark  amber-colored  bottle  in  which  a  small  quantity 
of  mercury  has  been  placed. 

Nitric  acid  test  solution,  HN03. — Add  36  mils  of  concentrated 
nitric  acid  to  290  mils  of  distilled  water  and  mix.  This  solution 
contains  approximately  10  per  cent,  of  absolute  HNO3. 

Oxalic  acid  test  solution,  H2C2O4  +  2H2O.— Dissolve  5  Gm.  of 
oxalic  acid  in  sufficient  distilled  water  to  measure  100  mils. 

Phenolphthalein  test  solution. — Dissolve  1  Gm.  of  phenol phthalein 
in  100  mils  of  alcohol.  It  gives  a  red  color  with  alkali  hydroxides 
or  carbonates,  and  acids  render  the  solution  colorless.  It  is  not 
suitable  as  an  indicator  for  ammonia  nor  in  the  presence  of  large 
quantities  of  ammonium  salts.  Phenolphthalein  paper  is  prepared 
by  impregnating  white,  unsized  paper  with  the  test  solution  and 
drying  it. 

Potassium  carbonate  test  solution,  I\2CO3. — Dissolve  10  Gm.  of 
potassium  carbonate  in  sufficient  distilled  water  to  measure  100  mils. 

Potassium  chloride  test  solution,  KC1. — Dissolve  10  Gm.  of  potas- 
sium chloride  in  sufficient  distilled  water  to  measure  100  mils. 

Potassium  chromate  test  solution,  K2CrQ4. — Dissolve  10  Gm.  of 
potassium  chromate  in  sufficient  distilled  water  to  measure  100  mils. 

Potassium  cyanide  test  solution,  KCN. — Dissolve  5  Gm.  of  potas- 
sium cyanide  in  sufficient  distilled  water  to  measure  100  mils. 

Potassium  dichromate  test  solution,  K2Cr2O?. — Dissolve  10  Gm. 
of  potassium  dichromate  in  sufficient  distilled  water  to  measure 
100  mils. 


SODIUM  HYPOBROMITE  TEST  SOLUTION  125 

Potassium  ferricyanide  test  solution,  K3Fe(CN)6.— Dissolve  1  part 
of  potassium  ferricyanide  in  about  10  parts  of  distilled  water.  This 
solution  should  be  freshly  made  when  required,  as  it  undergoes 
decomposition  with  formation  of  ferrocyanide  on  standing.  A 
freshly  prepared  aqueous  solution  mixed  with  ferric  chloride  T.S. 
which  has  been  well  diluted  with  distilled  water  shows  a  brown  tint, 
free  from  turbidity  or  a  shade  of  green. 

Potassium  ferrocyanide  test  solution,  I\4Fe(CN)6  +  3H20. — Dis- 
solve 10  Gin.  of  potassium  ferrocyanide  in  sufficient  distilled  water 
to  measure  100  mils. 

Potassium  hydroxide  test  solution,  KOH.— Use  the  official  solution 
of  potassium  hydroxide  containing  approximately  5  per  cent,  of 
KOH. 

Potassium  iodide  test  solution,  KI. — Dissolve  20  Gm.  of  potassium 
iodide  in  sufficient  distilled  water  to  measure  100  mils,  and  preserve 
the  solution  in  dark,  amber-colored,  well-stoppered  bottles.  The 
solution  should  be  frequently  renewed. 

Potassium  sulphate  test  solution,  K2£O4. — Dissolve  1  Gm.  of  potas- 
sium sulphate  in  sufficient  distilled  water  to  measure  100  mils. 

Potassium  sulphocyanate  test  solution,  KCNS.— Dissolve  1  Gm.  of 
potassium  sulphocyanate  in  sufficient  distilled  water  to  make  100 
mils. 

Silver  nitrate  test  solution,  AgNO3. — Dissolve  2  Gm.  of  silver 
nitrate  in  sufficient  distilled  water  to  measure  100  mils. 

Sodium  acetate  test  solution,  NaC2H3O2  +  3H2O.— Dissolve  10 
Gm.  of  sodium  acetate  in  sufficient  distilled  water  to  make  100  mils. 

Sodium  carbonate  test  solution,  Na2CO3  +  H2O. — Dissolve  10 
Gm.  of  monohydrated  sodium  carbonate  in  sufficient  distilled  water 
to  measure  100  mils. 

Sodium  cobaltic  nitrite  test  solution,  (NaNO2)6  Co2(NO2)6  +  H20.— 
Dissolve  4  Gm.  of  cobaltous  chloride  and  10  Gm.  of  sodium  nitrite 
in  about  50  mils  of  distilled  water,  add  2  mils  of  acetic  acid  and 
dilute  with  sufficient  distilled  water  to  measure  100  mils.  A  few 
drops  of  acetic  acid  should  be  added  to  the  solution  from  time  to 
time.  The  reagent  must  not  be  kept  longer  than  three  months. 
Should  any  precipitate  form  on  standing,  filter. 

Sodium  cyanide  test  solution,  NaCN. — Dissolve  1  Gm.  of  sodium 
cyanide  in  sufficient  distilled  water  to  measure  10  mils.  The  solu- 
tion must  be  freshly  prepared  when  required. 

Sodium  hydroxide  test  solution,  NaOH. — Dissolve  10  Gm.  of  sodium 
hydroxide  in  sufficient  distilled  water  to  measure  100  mils. 

Sodium  hypobromite  test  solution,  NaBrO.— To  a  solution  of  40 
Gm.  of  sodium  hydroxide  in  about  150  mils  of  distilled  water  add  10 
mils  of  bromine  and,  after  solution  has  taken  place,  add  sufficient 
distilled  water  to  measure  200  mils.  The  solution  must  be  freshly 
prepared  when  required  for  use. 


126  REAGENTS  AND  TEST  SOLUTIONS 

Sodium  nitroprusside  test  solution,  Na2FeNO(CN)5  +  2H2O.— 
Dissolve  1  part  of  sodium  nitroprusside  in  19  parts  of  distilled  water 
immediately  before  using. 

Sodium  phosphate  test  solution,  NaHPO4  +  12H2O. — Dissolve  10 
Gm.  of  sodium  phosphate,  in  clear  crystals,  in  sufficient  distilled  water 
to  measure  100  mils. 

Sodium  tartrate  test  solution,  Na2C4H4O6  +  2H2O.— Dissolve  10 
Gm.  of  sodium  tartrate  in  sufficient  distilled  water  to  measure  100 
mils. 

Sodium  thiosulphate  test  solution,  Na2S2O3  +  5H2O. — Dissolve  2.5 
Gm.  of  sodium  thiosulphate  in  sufficient  distilled  water  to  measure 
100  mils. 

Stannous  chloride  test  solution,  SnCl2  +  2H2O. — Dissolve  10  Gm. 
of  stannous  chloride  crystals  in  100  mils  of  distilled  water  to  which 
a  small  amount  of  hydrochloric  acid  has  been  added,  and  preserve 
the  solution  in  glass-stoppered  bottles  in  which  a  fragment  of  tin 
has  been  placed.  The  solution  must  be  renewed  at  frequent 
intervals. 

Stannous  chloride  test  solution,  saturated,  SnCl2  +  2H2O. — (For 
Bettendorf  s  test  for  arsenic.)  A  saturated  solution  of  stannous 
chloride  crystals  in  concentrated  hydrochloric  acid.  After  filtering 
the  solution  through  asbestos  it  has  not  more  than  a  pale  yellow 
color. 

Starch  test  solution.— Triturate  1  Gm.  of  cornstarch  with  10  mils 
of  cold  distilled  water,  add  boiling  distilled  water  with  constant 
stirring  to  make  about  200  mils,  then  boil  the  mixture  for  a  few 
minutes  until  a  thin,  translucent  fluid  is  obtained.  This  solution 
must  be  freshly  prepared  when  required. 

Sulphuric  acid  test  solution,  H2S04. — Use  the  official  diluted  sul- 
phuric acid,  containing  about  10  per  cent,  of  absolute  H2SO4. 

Tartaric  acid  test  solution,  H2C4H4O6. — Dissolve  1  part  of  tartaric 
acid  in  3  parts  of  distilled  water.  This  solution  must  be  frequently 
renewed. 

Turmeric  tincture — Digest  any  convenient  quantity  of  ground 
turmeric  root  repeatedly  with  small  quantities  of  distilled  water  and 
discard  the  liquids.  Then  digest  the  dried  residue  for  several  days 
with  six  times  its  weight  of  alcohol,  and  filter. 

Turmeric  paper. — Impregnate  white,  unsized  paper  with  the 
tincture,  and  dry  it.  The  tincture,  as  well  as  the  paper,  turns  brown 
with  alkalies,  and  the  original  yellow  color  is  restored  by  acids,  with 
the  exception  of  boric  acid,  which,  especially  in  the  presence  of 
hydrochloric  acid,  turns  the  color  to  reddish-brown,  which  is  changed 
to  bluish-black  by  ammonia. 

Zinc  sulphate  test  solution,  ZnSO4. — Dissolve  10  Gm.  of  zinc  sul- 
phate in  sufficient  distilled  water  to  measure  100  mils. 


INDEX 


ACETIC  acid,  description  of,  101 

reagent,  120 

salts,  101-102 

tests,  102 
Acid,  acetic,  101 
anhydrid,  16 
arsenic,  39 
boric,  101 
bromauric,  38 
chlorauric,  38 
chloroplatinic,  38 
dibasic,  15 
hydr-,  16 
hydriodic,  89 
hydrobromic,  88 
hydrochloric,  86 

test  solution,  122 
hydrocyanic,  90 
monobasic,  15 
muriatic,  86 
nitric,  102 

test  solution,  124 
orthpphosphoric,  98 
oxalic,  97 
oxy-,  16 

phosphoric,  98-99 
radicals,  16 
reaction,  16 
sulphuric,  93-95 

test  solution,  126 
sulphurous,  95 
tetra-basic,  15 
tri-basic,  15 
Acids,  classification  of,  86 

group  A,  86-93 

group  B,  93-99 

group  C,  100-103 
description  of,  15 
Adhesion,  11 
Alkali,  16 

Alkaline  reaction,  16 
Alloy,  15 
Alum,  52 
burnt,  52 
dried,  52 

examination  of,  114 
exsiccated,  52 


Alum,  ferric,  48,  94 

Aluminum  and  ammonium  sulphate, 

52,  94 

and  potassium  sulphate,  52,  94 
and  sodium  sulphate,  52 
chloride,  52 
compounds,  52 
description  of,  52 
hydroxide,  52 
oxide,  52 
silicates,  52 
sulphate,  52,  94 
tests,  52-53 
Amalgams,  25 
Ammonia,  80 
Ammoniated  mercury,  29 
Ammonium  acetate,  80,  101 
alum,  52,  94 

description  of,  114 
tests  for  dentity  of,  111 
for  impurities  of,  114 
bicarbonate,  96 
bromide,  80,  88 
description  of,  111 
tests  for  identity  of,  111 
for  impurities  of,  111 
carbamate,  96 
carbonate,  80,  96 

test  solution,  120 
chloride,  80,  87 

test  solution,  120 
compounds,  80 
description  of,  79 
hydroxide,  80 

test  solution,  120 
hypophosphite,  80 
iodide,  80,  89 

molvbdate  test  solution,  120 
nitrate,  80,  103 
oxalate,  80,  98 

test  solution,  120 
phosphate,  80 
polysulphide,  80,  91 
test  solution,  121 
sulphate,  80,  94 
sulphide,  80,  91 
tests,  80 

Amorphous  compounds,  20 
Analysis,  chemical,  21 


128 


INDEX 


Analysis,  qualitative,  21 

quantitative,  21 
Analytical  tables,  35 
Anhydrous  calcium  sulphate,  67,  94 
Antimonous  chloride,  42 

oxide,  42 

sulphide,- 42,  91 
Antimony  and  potassium  tartrate,  42 

compounds,  42 

test  for  arsenic  in,  117 

description  of,  41-4b 

tests,  42-43 

Antimonyl-potassium  tartrate,  42 
Apparatus  for  arsenic  test,  115 
Aqua  fortis,  102 
Argenti  nitras,  26 

oxidum,  26 
Argentum,  26 
Arsenic  acid,  39 

antidote,  48 

compounds,  39 

description  of,  39 

tests,  39-41,  115-117 

trioxide,  39 

white,  39 

Arseniuretted  hydrogen,  39 
Arsenous  anhydride,  39 

iodide,  39,  89 

oxide,  39 

sulphide,  39 
Arsine,  39 
Asbestos,  75 
Atomic  theory,  1 1 

weight,  14 

Atoms,  definition  of,  11 
Auri  et  sodii  chloridum,  38 
Aurum,  37 


B 


BAKING  soda,  78,  96 
Barium  carbonate,  69 

chloride,  69,  87 
test  solution,  121 

compounds,  69-70 

description  of,  69 

dioxide,  69 

hydroxide,  70 
test  solution,  121 

nitrate,  70,  103 
test  solution,  121 

peroxide,  69 

sulphate,  70,  94 

sulphide,  70,  91 

tests,  70 
Bases,  15 
Basic  copper  acetate,  32,  102 

zinc  carbonate,  61 
Bettendorf's  arsenic  test,  117 
Binary  compound,  17 
Biniodide  of  mercury,  29,  89 
Bismuth  chloride,  30 


Bismuth,  compounds,  30 

test  for  arsenic  in,  117 
description  of,  30 
nitrate,  30,  103 
oxide,  30 
subcarbonate,  30 
subnitrate,  30,  103 
description  of,  113 
tests  for  identity,  113 

for  impurities",  113-114 
tests,  31 
Bismuthyl,  30 
Bisulphite  of  lime,  95 
Blank  test,  40 
Bleaching  powder,  67 
Blue  stone,  94 

vitriol,  32,  94 
Bond,  14 
Boracic  acid,  100 
Borax,  78,  101 

Boric  acid,  description  of,  100 
salts,  101 
tests,  101 
Boroglycerin,  101 
Bromauric  acid,  38 
Bromine  test  solution,  121 

water,  121 
Burnt  alum,  52 
Butter  of  antimony,  42 


CADMIUM  bromide,  33 

chloride,  33 

compounds,  33 

description  of,  33 

iodide,  33 

nitrate,  33 

oxide,  33 

sulphate,  33,  94 

sulphide,  33 

tests,  33-34 
Calamine,  96 
Calcined  magnesia,  75 
Calcium  acetate,  67 

bisulphite,  95 

bromide,  67,  68 

carbide,  67 

carbonate,  67,  96 

chloride,  67,  87 
test  solution,  121 

compound,  67 

description  of,  67 

hydroxide,  67 
test  solution,  121 

hypophosphite,  67 

iodide,  67 

oxalate,  98 

oxide,  67 

phosphate,  67 

sulphate,  67,  94 


INDEX 


129 


Calcium  sulphate  test  solution,  121 

sulphide,  67,  91 

sulphite,  67,  95 

tests,  67-68 
Calomel,  27,  87 
Caustic  potash,  76 
test  solution,  125 

soda,  78 

test  solution,  125 
Cerium  oxalate,  98 
Cerussa,  23 
Chalk,  67,  96 
Changes,  chemical,  9 

physical,  9 
Chemical  affinity,  11 

analysis,  21 

attraction,  11 

changes,  9,  12 

equations,  13 

formulas,  13 

reaction,  10 
Chemism,  11 
Chemistry  denned,  10 
Chili  saltpetre,  78,  103 
Chlorauric  acid,  38 
Chloride  of  lime,  67 
Chlorinated  lime,  67 
Chlorine  test  solution,  121 

water,  121 

Chloroplatinic  acid,  38 
Chrome  alum,  51,  94 

yellow,  23,  51 
Chromic  acid,  51 

anhydride,  51 

oxide,  51 
Chromium  compounds,  51 

description  of,  51 

potassium  sulphate,  51,  94 

sulphate,  51 

tests,  51 

trioxide,  51 
Cinnabar,  29,  92 
Cobalt  chloride,  57 

compounds,  57 

description  of,  57 

nitrate,  57 

oxide,  57 

sulphate,  57,  94 

tests,  58 

Cobaltous  chloride  test  solution,  121 
Cohesion,  11 
Colloids,  20 

Colorless  ammonium  sulphide,  91 
Combination,  12 
Composition,  9 
Compound  denned,  11 

radicals,  16 
Compounds,  aluminum,  52 

ammonium,  79 

amorphous,  20 

anhydrous,  21 

antimony,  42 
9 


Compounds,  arsenic,  39 

barium,  69 

binary,  17 

bismuth,  30 

cadmium,  33 

calcium,  67 

chromium,  51 

cobalt,  57 

copper,  32 

crystalline,  20 

efflorescent,  21 

ferric,  48 

ferrous,  48 

gold,  38 

lead,  23-24 

lithium,  79 

magnesium,  74—76 

manganese,  60 

mercuric,  29 

mercurous,  25 

nickel,  59 

platinum,  38 

potassium,  76-77 

silver,  26 

sodium,  77-79 

stannic,  43-44 

stannous,  43^14 

strontium,  68 

tin,  43-44 

zinc,  61 

Copper  acetate,  101 
basic,  32,  102 

aceto-arsenite,  32,  39,  101 

arsenite,  32 

chloride,  32 

compounds,  32 

description  of,  31-33 

nitrate,  32 

subacetate,  102 

sulphate,  32,  94 
test  solution,  121 

tests  for,  32-33 
Copperas,  48,  94 
Corrosive  mercuric  chloride,  87 

sublimate,  29 
Cream  of  tartar,  76 
Crystalline  compound,  20 
Crystallization,  water  of,  20 
Crystalloids,  20 
Crystals,  20 
Cupri  sulphas,  32 
Cupric  chloride,  32 

nitrate,  32 

oxide,  32 

sulphate,  32,  94 

test  solution,  121 
Cuprous  oxide,  32 
uprum,  31 


DECOMPOSITION,  12 
double,  12 


130 


INDEX 


Definite  proportions,  law  of,  10 

Deliquescent  substances,  21 

Destruction  of  organic  matter,  119 

Dialysis,  20 

Di-basic  acid,  15 

Diffusion,  20 

Diluted  acetic  acid,  101 

hydriodic  acid,  89 

hydrobromic  acid,  88 

hydrochloric  acid,  86 
description  of,  1 12 
tests  for  identity,  112 
for  impurities,  112 

hydrocyanic  acid,  90 

phosphoric  acid,  98 

sulphuric  acid,  93 
Di-sodium  hydrogen  phosphate,  78 

test  solution,  126 
Dissociation,  20 
Divisibility,  9 
Dried  alum,  52 
Dyad,  14 


E 


EFFLORESCENT  compound,  21 
Electrolysis,  20 
Element,  11 

characteristic,  15 

monad,  14 
Energy,  conservation  of,  10 

definition  of,  10 
Epsom  salt,  75,  94 
Equations,  chemical,  13 
Ethiops  mineral,  29 
Exsiccated  alum,  52 

calcium  sulphate,  67 

ferrous  sulphate,  48 

sodium  arsenate,  77 
phosphate,  78 
sulphite,  78,  95 
Exsiccation,  21 
Extension  of  matter,  9 


F 


FERRIC  acetate,  102 
alum,  48,  94 

ammonium  sulphate,  48,  94 
chloride,  48,  87 

test  solution,  121 
ferrocyanide,  48 
hydroxide,  48 
hypophosphite,  48 
nitrate,  48,  103 
oxide,  48 
phosphate,  48 
pyrophosphate,  48 
subsulph.ate,  48,  94 
sulphate,  48,  94 


Ferroso-ferric  oxide,  48 
Ferrous  bromide,  48 
carbonate,  48,  96 
chloride,  48 
hydroxide,  48 
iodide,  48,  89 
oxide,  48 
sulphate,  48,  94 
description  of,  110 
test  solution,  121 
tests  for  identity,  110 
for  impurities,  111 
sulphide,  48,  92 
Filtrate,  19 
Formulas,  13 


GALENA,  24,  92 
Glacial  acetic  acid,  101 
Glass,  soluble,  78 

water,  78 

Glauber's  salt,  78,  94 
Glyceryl  borate,  101 
Gold  and  sodium  chloride,  38 

chloride,  38,  87 

compounds,  38 

description  of,  37 

tests,  38 
Gravitation,  9 
Green  iodide  of  mercury,  27 

vitriol,  48,  94 
Group  reagents,  23 
Gypsum,  67,  94 


HEAVY  magnesia,  75 
magnesium  oxide,  75 
metals,  test  for,  118 
spar,  94 
Hematite,  48 
Hexad,  14 
Hydracid,  16 
Hydrargyri  chloridum  mite,  25 

corrosivum,  29 
iodidum  flavum,  25 

rubrum,  29 
oxidum  flavum,  29 

rubrum,  29 
Hydrargyrum,  25,  29 

ammoniatum,  29 
Hydriodic  acid,  description  of,  89 
salts,  89 
tests,  89-90 

Hydrobromic  acid,  description  of,  88 
salts,  88 
tests,  88 

Hydrochloric  acid,  description  of,  86 
salts,  87 


INDEX 


131 


Hydrochloric  acid  test  solution,  122 

tests,  87 
Hydrocyanic  acid,  description  of,  90 

salts,  90 

tests,  90-91 

Hydrogen  antimonide,  42 
arsenide,  39 
phosphide,  117 
sulphide,  91 

test  solution,  122 
Hydrosulphuric  acid,  description  of,  91 

salts,  91-92 

tests,  92 

Hygroscopic  substance,  21 
Hyposulphite  of  soda,  78 

test  solution,  126 


IDENTITY  tests,  110 

Impenetrability,  9 

Impurities,  109 

Indestructibility,  9 

Indicators,  16 

Inertia,  9 

Inorganic  chemicals,  qualitative  exami- 
nation of  official,  109-114 

Iodine  test  solution,  122 

Ions,  20 

Iron  compounds,  48 
description  of,  48 
tersulphate,  97  . 
tests,  49-50 


KNOWN  solutions,  23 


LAW  of  definite  proportions,  10 

of  multiple  proportions,  10 

of  precipitation,  19 
Lead  acetate,  23,  102 

carbonate,  23 

chromate,  23,  51 

compounds,  23-24 

description  of,  23 

dioxide,  24 

glass  wool,  122 

iodide,  24,  89 

nitrate,  24,  103 

oxide,  24 
red,  24 

peroxide,  24 

sub-acetate,  102 

sub-carbonate,  96 

sulphide,  24,  92 

test  paper,  122 


Lead  test  solution,  122 

tests,  24 
Lime,  67 

chloride  of,  67 

quick,  67 

slaked,  67 

water,  67,  121 
Limestone,  67,  96 
Litharge,  24 
Lithium  bromide,  79,  88 

carbonate,  79,  96 

chloride,  79 

compounds,  79 

description  of,  79 

tests,  79 
Litmus  paper,  preparation  of,  112-223 

test  solution,  122-123 
Liver  of  sulpher,  77 
Lunar  caustic,  26,  103 
test  solution,  125 


M 

MAGNESIA,  75,  96 

alba,  75 

mixture,  123 
Magnesite,  75,  96 
Magnesium  carbonate,  75,  96 

chloride,  75,  87 

compounds,  75 

description  of,  74-75 

hydroxide,  75 

oxide,  75 

silicates,  75 

sub-carbonate,  75 

sulphate,  75,  94 
test  solution,  123 

tests,  75 
Magnetite,  48 
Manganese  chloride,  60 

compounds,  60 

description  of,  59-60 

dioxide,  60 

hypophosphite,  60 

sulphate,  60 

tests,  60-61 
Manganous  chloride,  60 

hypophosphite,  60 

sulphate,  60,  94 

test  solution,  123 
Marble,  67,  96 

Marsh's  test  for  arsenic,  40-41 
Mass  defined,  11 
Matter  changes,  9 

composition  of,  11 

conservation  of,  10 

definition  of,  9 

structure  of,  11 
Mayer's  reagent,  124 
Meerschaum,  75 
Mercuric  bromide,  29 


132 


INDEX 


Mercuric  bromide  test  paper,  123 
solution,  123 

chloride,  29,  87 
test  solution,  123 

cyanide,  29,  90 

iodide,  29,  89 

mercury  compounds,  29 
description  of,  29 
tests,  29-31 

nitrate,  29,  103 

oxide,  red,  29 
yellow,  29 

potassium  iodide  test  solution,  124 

subsulphate,  29 

sulphate,  29,  94 

sulphide,  black,  29 

red,  29,  92 

Mercuric-ammonium  chloride,  29 
Mercurous  chloride,  25,  87 

iodide,  25,  89 

mercury  compounds,  25 
description  of,  25 
tests,  25-27 

nitrate,  25 

test  solution,  124 

oxide,  25 

sulphate,  94 
Metal  defined,  15 
Metals,  classification  of,  22 

Group  1,  22,  23-28 

Group  2,  22,  28-35 

Group  3,  22,  37-47 

Group  4,  22,  48-54 

Group  5,  22,  57-63 

Group  6,  22,  67-71 

Group  7,  22,  74-81 
Mild  mercurous  chloride,  87 
Minium,  24 
Miscible  liquids,  19 
Mixture  defined,  12 
Molecules,  11 
Monad,  14 
Monobasic  acid,  15 
Monohydrated  sodium  carbonate,  78, 

96 

Monsel's  salt,  48,  94 
Mosaic  gold,  43 

Multiple  proportion,  law  of,  10 
Muriatic  acid,  86 


N 


NASCENT  state,  12 
Nessler's  reagent,  124 
Neutralization,  15 
Nickel  compounds,  59 

description  of,  58-£9 

tests,  59 
Nickel-ammonium  chloride,  59 

oxide,  59 

sulphate,  59,  94 


Nitre,  76,  103 

Nitric  acid,  description  of,  102 

salts,  103 

test  solution,  124 

tests,  103 
Nomenclature,  17 
Normal  salts,  16 


OIL  of  vitriol,  93 

Organic  matter,  destruction  of,  119 

Orpiment,  39 

Ortho-arsenic  acid,  39 

Orthoboric  acid,  100 

Osrrosis,  20 

Oxalic  acid,  description  of,  97-98 

salts,  98 

test  solution,  1,24 

tests,  98 
Oxyacids,  definition  of,  16 


PARIS  green,  32,  39,  101 
Pearlash,  76 
Pentad,  14 

Permanent  white,  70,  94 
Phenolphthalein  test  solution,  124 
Phosphoric  acid,  description  of,  98 
salts,  99 
tests,  99 
Physical  changes,  9 

science,  10 
Physics,  10 

Plaster  of  Paris,  67,  94 
Platinum  chloride,  38,  87 
compounds,  38 
description  of,  38 
tests,  39 
Plumbi  acetas,  23 

oxidum,  24 
Plumbum,  23 
Pot  ash,  76 
alum,  94 
caustic,  76 

test  solution,  125 
prussiate  of,  76 
red,  76 
yellow,  76 
Potassa,  76 

test  solution,  125 
Potassium  acetate,  76,  102 
acid  oxalate,  98 
alum,  52 

description  of,  114 
tests  for  identity,  114 
for  impurities,  114 
and  sodium  tartrate,  77 


INDEX 


133 


Potassium  arsenite,  39 
bicarbonate,  76,  96 
bitartrate,  76 
bromate,  76 
bromide,  76,  88 
carbonate,  76,  96 

test  solution,  124 
chlorate,  76 
chloride,  76 

test  solution,  124 
chromate,  51,  76 

test  solution,  124 
compounds,  76-77 
cyanide,  76,  90 
*  test  solution,  1 24 
description  of,  76 
dichromate,  51,  76 
test  solution,  124 
ferri cyanide,  76 

test  solution,  125 
hydrate,  76 

test  solution,  125 
hydroxide,  76 

test  solution,  125 
hypophosphite,  76 
iodide,  76,  89 

test  solution,  125 
manganate,  60 
nitrate,  76,  103 
perchlorate,  76 
permanganate,  60,'  76 
sulphate,  76,  94 

test  solution,  125 
sulphocyanate,  77 

test  solution,  125 
tests,  77 
thiocyanate,  77 

test  solution,  125 
Precipitate,  19 
red,  29 
white,  29 
Precipitated    calcium    carbonate,    67, 

96 

phosphate,  67 
manganese  dioxide,  60 
zinc  carbonate,  61 
Preparation  of  chemicals  to  be  tested 

for  arsenic,  1 16 
Prepared  chalk,  67,  96 
Proportions,  definite,  10 

law  of,  10 

Proto-iodide  of  mercury,  27,  89 
Prussian  blue,  48 
Prussiate  of  potash,  76 
red,  76 
yellow,  76 
Prussic  acid,  90 
Purity,  109 
chemical,  109 
technical,  109 
tests,  110 
Pyrolusite,  60 


QUALITATIVE  analysis  defined,  21 
examination    of     official    inorganic 

chemicals,  109-114 
tests  of  the  U.  S.  P.,  general,  115 

Quantitative  analysis  defined,  21 

Quicklime,  67 

Quicksilver,  25 


RADICALS,  acid,  16 

compounds,  16 
Reaction,  acid,  16 

alkaline,  16 

characteristic,  21 

neutral,  16 
Reagents,  21,  120 

group,  23 

Mayer's  124 

Nessler's  124 
Red  lead,  24 

mercuric  iodide,  89 
oxide,  24 

precipitate,  29 

prussiate  of  potash,  76 

test  solution,  125 
Reinsch's  test  for  arsenic,  40 
Rochelle  salt,  77 


S 


SAL  ammoniac,  80,  87 

soda,  78 

volatile,  80 
Saleratus,  76,  96 
Salt,  78,  87 

Glauber's,  78,  94 

of  lemon,  98 

of  tartar,  76 

rochelle,  77 

spirit  of,  86 
Saltpetre,  76,  103 

Chili,  78,  103 
Salts,  15 

acid,  16 

basic,  16 

double,  16 

normal,  16 

of  acetic  acid,  101 

of  boric  acid,  101 

of  carbonic  acid,  96 

of  hydriodic  acid,  89 

of  hydrobromic  acid,  88 

of  hydrochloric  acid,  87 

of  hydrocyanic  acid,  90 

of  hydrosulphuric  acid,  91-92 

of  nitric,  acid,  103 

of  oxalic  acid,  98 

of  phosphoric  acid,  99 


134 


INDEX 


Salts  of  sulphuric  acid,  94 

of  sulphurous  acid,  95 
Samples : 

No  1,  analysis  for  Group  1  of  metals, 

27-29 
No.  2,  analysis  for  Group  2  of  metals, 

34-35 
No.  3,  analysis  for  Group  1,  2  of 

metals,  36-37 
No.  4,  analysis  for  Group  3  of  metals, 

44-46 
No.  5,  analysis  for  Group  1,  2,  3  of 

metals,  46-47 
No.  6,  analysis  for  Group  4  of  metals, 

53-54 
No.  7,  analysis  for  Groups  1,  2,  3,  4 

of  metals,  55-57 
No.  8,  analysis  for  Group  5  of  metals, 

63 
No.  9,  analysis  for  Groups  1,  2,  3,  4, 

5  of  metals,  64-66 
No.    10,   analysis   for   Group   6   of 

metals,  70-71 
No.  11,  analysis  for  Groups  1-6  of 

metals,  71-74 
No.  12,  analysis    for   -Group    7    of 

metals,  81 
No.   13,  analysis  for  all  groups  of 

metals,  82-85 
No.    14,   analysis   for   Group   A   of 

acids,  92-93 
No.  15,  analysis  for  important  metals 

and    for    radicals    of    sulphuric, 

sulphurous  and  carbonic  acids,  97 
No.  16,  analysis  for  principal  metals 

and  for  radicals  of  oxalic  and  phos- 
phoric acids,  99-100 
No.  17,  analysis  for  principal  metals 

and  for  radicals  of  boric,  acetic 

and  nitric  acids,  103-104 
No.  18,  analysis  for  principal  metals 

and  acids,  104-105 
Nos.  19  and  20,  qualitative  analysis 

of  a  solid,  106-108 
No.  21,  examination  of    granulated 

ferrous  sulphate  for  identity  and 

purity,  110^111 
No.  22,  examination  of  ammonium 

bromide  for  identity  and  purity, 

111-112 

No.  23,  examination  of  diluted  hydro- 
chloric acid  for  identity  and  purity, 

112 

No.  24,  examination  of  bismuth  sub- 
nitrate  for  identity   and   purity, 

113-114 
No.    25,    examination   of  alum  for 

identity  and  purity,  114 
Saturated  solution,  19 
Scheele's  green,  32 
Sediment,  19 
Silver  bromide,  88 


Silver  chloride,  26 
compounds,  26 
cyanide,  26,  90 
description  of,  26 
iodide,  26,  89 
nitrate,  26,  103 

test  solution,  125 
oxide,  26       . 
sulphate,  26,  94 
tests,  26-27 
Slaked  lime,  67 

test  solution,  121 
Smalt,  57 
Soapstone,  75 
Soda,  78 

baking,  78,  96 
caustic,  78 

test  solution,  125 
hyposulphite  of,  78 
sal,  78 

washing,  78,  96 
Sodium  acetate,  77,  102 

test  solution,  125 
alum,  52 
arsenate,  39,  77 
bicarbonate,  78,  96 
bisulphite,  78,  95 
borate,  78,  101 
bromide,  78,  88 
carbonate,  78,  96 

test  solution,  125 
chlorate,  78 
chloride,  78,  87 

cobaltic  nitrite  test  solution,  125 
compound,  77-78 
cyanide,  78,  90 

test  solution,  125 
description  of,  77 
di-hydrogen  phosphate,  78 
hydrate,  78 

test  solution,  125 
hydroxide,  78 

test  solution,  125 
hypobromite  test  solution,  125 
hypochlorite,  78 
hypophosphite,  78 
iodide,  78,  89 
nitrate,  78,  103 
nitrite,  78 

nitroprusside  test  solution,  126 
oxalate,  78,  98 
perborate,  78 
peroxide,  78 
phosphate,  78 

test  solution,  126 
pyroborate,  78,  101 
pyrophosphate,  78 
silicates,  78 
stannate,  43 
sulphate,  78,  94 
sulphide,  78,  92 
sulphite,  78,  95 


INDEX 


135 


Sodium  sulphite,  exsiccated,  78,  95 

tartrate,  78 

test  solution,  126 

tetraborate,  78,  10 1 

thiosulphate,  78 
test  solution,  126 

tests,  78-79 
Solids,  analysis  of,  106 
Soluble  glass,  78 
Solution,  19 

known,  23 

saturated,  19 

super-saturated,  19 

test,  21,  115 

unknown,  23 
Solvent,  19 
Specific  properties,  9 
Spirit  of  salt,  86 
Standard  arsenic  test  solution,  115 

stain,  116 
Stannic  chloride,  43,  87 

compounds  43-44 

sulphide,  43 

Stannous  chloride,  43,  87 
test  solution,  126 

compounds,  43-44 

sulphide,  43 
Stannum,  43 
Starch  test  solution,  126 
Stibine,  42 
Stibium,  41 
Strontium  bromide,  68,  88 

carbonate,  68 

chloride,  68 

compounds,  68 

description  of,  68 

iodide,  68,  89 

nitrate,  68,  103 

sulphate,  98 
Substance,  composition  of,  9 

definition  of,  9 

deliquescent,  21 

hygroscopic,  21 

neutral,  16 

specific  properties  of,  9 
Substitution,  12 
Sugar  of  lead,  23,  102 
Sulphuretted  hydrogen,  91 

potash,  77 

Sulphuric  acid,  description  of,  93 
salts,  94 

test  solution,  126 
tests,  94-95 

Sulphurous  acid,  description  of,  95 
salts,  95 
tests,  95-96 

Super-saturated  solution,  19 
Symbols,  13 


TABLES,  analytical,  description  of,35-36 
metals,  groups  1-2,  36-37 


Tables,  metals,  groups,  1-3,  46-40 
1-4,  55-57 
1-5,  64-66 
1-6,  71-74 
1-7,  82-85 
Talcum,  75 
Tartar,  cream  of,  76 
emetic,  42 
salt  of,  76 

Tartaric  acid  test  solution,  126 
Test  apparatus  for  arsenic,  115 

solution,  21,  115,  120 
Tests,  special,  Bettendorf's,  117 
Marsh's  40-41 
Reinsch's,  40 
Tetra-basic  acid,  15 
Tetrad,  14 
Tin  compounds,  43 
description  of,  43 
tests  for,  43-44 
Triad,  14 
Tri-basic  acid,  15 
Tri-sodium  phosphate,  78 
Turmeric  paper,  126 

tincture,  126 
Turpeth  mineral,  29 


UNKNOWN  solutions,  23 


VALENCE,  14 
Verdigris,  32,  102 
Vermilion,  29,  92 
Vitriol,  blue,  32,  94 
green,  48,  94 
oil  of,  93 
white,  61,  94 


WASHING  soda,  78,  96 
Water,  bromine,  121 

chlorine,  121 

glass,  78 

lime,  121 

of  crystallization,  20 
Weight,  atomic,  14 

combining,  14 

molecular,  14 
White  arsenic,  39 

lead,  23,  96 

permanent,  70 

precipitate,  29 

vitriol,  61,  94 
Whiting,  67 


136 


INDEX 


YELLOW  ammonium  sulphide,  80,  91 

test  solution,  121 
mercuric  oxide,  29 
mercurous  iodide,  89 
prussiate  of  potash,  76 

test  solution,  125 


Z 

ZINC  acetate,  61,  102 


Zinc  bromide,  61 
carbonate,  96 
chloride,  61,  87 
compounds,  61 
description  of,  61 
iodide,  61,  89 
oxide,  61 

sub-carbonate,  61 
sulphate,  61,  94 

test  solution,  126 
sulphide,  92 
tests,  62 
white,  61 


QD83- 

B81 

1920 


T 


